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TEMPLE / Science / MLSC 1031 / What are the main states of matter?

What are the main states of matter?

What are the main states of matter?


School: Temple University
Department: Science
Course: General Chemistry I
Professor: Bloxton
Term: Spring 2016
Cost: 50
Description: These notes were prepared using the study guides for all of the previous four exams and well as the final exam study guide, and also include the topics which will be on the final but were not covered in the previous exam. They are detailed and provide examples.
Uploaded: 12/12/2017
57 Pages 66 Views 2 Unlocks


What are the main states of matter?


•Topic 1: States of Matter

∙ Matter is anything that has mass and volume

∙ 3 states

∙ Solid

 Crystalline

∙ Atoms or molecules are in a regular pattern and fixed positions ∙ Atoms or molecules vibrate in place and don’t move

∙ Ex: diamond  

 Amorphous

∙ no regular pattern

∙ random arrangement of atoms and molecules

∙ ex: charcoal

∙ Liquid

 Atoms/ molecules not packed together tightly

What is the meaning of crystalline in the matter?

 Atoms/molecules free to move about

 Fixed volume

 Atoms/molecules assume shape of the container

∙ Gas

 Compression possible because lots of space between  


 Random movement of atoms/molecules with high velocity

 Atoms/molecules assume shape and volume of the container •Topic 2: Types of Mixtures

∙ Homogenous

∙ Separable by distillation

∙ Heterogeneous

∙ Separable by filtration  

•Topic 4: Physical vs Chemical Properties  We also discuss several other topics like How long do i have to work for a company before i get maternity pay?

∙ Physical

What is the meaning of amorphous in the matter?

∙ Properties which a substance displays without changing its chemical  composition

∙ Properties which do not depend on the amount of substance 1


∙ Ex: MP, BP, density, color, smell

∙ Chemical

∙ Properties which a substance displays only by changing its composition  during a chemical change

∙ Ex: flammability, corrosiveness, acidity, alkalinity, toxicity

•Topic 5: Physical vs Chemical Changes

∙ Physical

∙ Change in the state of matter or appearance

 Sublimation: solid to gas

 Deposition: gas to solid

∙ No change in chemical identity

∙ Changes require input or removal of energy  

∙ Chemical

∙ Changes that alter the composition of matter

•Topic 6: Basic SI Units and Prefix Multipliers

SI Base Units
















Amount of substance



Electric current



Luminous intensity



We also discuss several other topics like In what two separate ways do plants obtain nutrients?

SI Prefix Multipliers













If you want to learn more check out When did the nubians conquer egypt?
If you want to learn more check out How long did it take for the buildings to collapse?
























If you want to learn more check out How can i improve my careless mistakes?

•Topic 7: Dimensional Analysis involving Mass, Volume and  Density

∙ Dimensional analysis – using units as a guide to solving problems  ∙ Use the formula

Information given x conversion factor(s) = information sought

Given unit x desired unit/given unit = desired unit

•Topic 8: Significant Figures

∙ The greater the number of significant figures, the greater the certainty of the  measurement

∙ Significant figures = those which are the non-lace-holding digits ∙ To determine the number of significant figures

∙ We must distinguish between zeroes that are significant and those that  simply mark the decimal place

∙ The leading zeroes mark the decimal place, so are not significant ∙ Trailing zeroes after a decimal point are significant  

∙ Trailing zeroes before a decimal always significant If you want to learn more check out When did john locke write “an essay concerning human understanding”?

∙ Trailing zeroes before an implied decimal points are ambiguous ∙ All nonzero digits are significant

∙ Interior digits are significant

∙ How to multiply and divide  

∙ How to subtract and add

•Topic 9: Scientific Notation

∙ Used to avoid ambiguity in significant figures

∙ How to multiply and divide



∙ How to subtract and add


•Topic 1: Laws of Basic Chemistry  

∙ Law of conservation of mass (Lavoisier)

∙ Law of definite proportion – all samples of a given compound have the same  proportions as their constituent elements

∙ Law of multiple proportion – when two elements (A and B) form two different  compounds, the mass of B that combines with 1 g of A can be expressed as a  ratio of small whole numbers

•Topic 2: Atomic Theory  

∙ Atoms first seen through scanning tunneling microscopy

∙ First thought by Leucippus & Democritos that atoms were uncuttable and that  nothing exists except atoms and empty space

∙ Dalton’s Atomic Theory

∙ All matter is composed of atoms

∙ Atoms cannot be made or destroyed

∙ All atoms of same elements are identical

∙ Different elements have different atoms

∙ Chemical reactions occur when atoms rearrange

∙ Compound form from atoms of constituent elements

∙ Thomson’s discovery of the electron

∙ Experiment: cathode ray tube

∙ Cathode rays are a stream of negatively charged particle

∙ Deduced charge-to-mass ratio: 1.76X108 C/ 1 g  

∙ Electron is 2000X lighter than a H atom

∙ Significance: proved that atoms could be further divided  

∙ Millihan’s discovery of the charge of the electron

∙ Oil drop experiment

∙ Negatively charged drops suspended between two plates

∙ Using the voltage and mass of the drop, Millihan found the charge in each  drop

∙ Always a multiple of -1.60X10-19 C

∙ Therefore, charge of one electron is -1.60 X 10-19 C



∙ Given the mass-to-charge ratio and charge of a single electron, the mass  of the electron is 9.1 X 10-28 g

∙ Significance: the atoms is neutral and electrons are negatively charged, so there must be a positively charged particle in the atom

∙ Rutherford’s atomic model

∙ Proposed a positively charged nucleus

∙ Electrons surround the nucleus

∙ The rest of the atom is empty space

∙ Plum pudding model

∙ Experiment: gold foil experiment  

∙ Alpha particles aimed at gold foil

∙ As many negative particles as positive particles

∙ Problem: Something must account for extra mass, if not protons ∙ Then, Chadwick proposed neutrons with same mass as proton •Topic 3: Subatomic particles  

∙ Charges, Weights, Symbols


Symbo l

Relativ e  










+1.60 X  


1.6762 x  

10-27 kg

1.00 amu






1.67493 x  10-27 kg

1.00 amu





-1.60 X 10- 19

9.10 x 10-31  kg

5.48 x 10- 4 amu

∙ 1 a.m.u. = 1/12 (mass of an atom of carbon-12) = 1.66 x 10-27 kg ∙ Protons and atomic number  

∙ Atomic number (Z) is unique to each element

∙ Z = # of protons

∙ Z, and therefore # of protons, defines the element

∙ Different elements contain different numbers of protons

∙ Each element has a name and a symbol  

•Topic 4: Isotopes



∙ Dalton proposed that all atoms of a given elements have the same mass, but  this is not true!

∙ Isotope = atoms that contain same number of p+ but different of n0 ∙ Mass number (A) = #n0 + #p+  

∙ Percent abundance calculations

∙ Mass of an isotope is measured through mass spectroscopy

∙ Atomic masses are weighted according to the natural abundance of each  isotope

∙ Atomic mass = [(percent abundance of isotope 1 * weight of isotope 1) +  (percent abundance weight of isotope 2 * weight of  

isotope 2)]/100

•Topic 5: Periodic law

∙ Periodic Law – elements with similar properties occur with a similar patten  when organized in order of increasing mass

∙ Most elements are metals

∙ Non-metals on right

∙ Metalloids: B, Si, Ge, As, Sb, Te, At

∙ Group Names

∙ Group 1 – Alkaline metals; produce alkaline solution in water

∙ Group 2 – alkaline earth metals

∙ Group 15 – Pnictogens

∙ Group 16 – Chalcogens

∙ Group 17 – Halogens

∙ Group 18 – Noble gases

∙ Period

•Topic 6: Ions

∙ Ions form when atoms lose or gain electrons

∙ To represent charge, remember that the number always goes before the  charge

∙ How to predict charges

∙ Metals tend to form cations by losing electrons  

∙ Nonmetals tend to gain electrons and form anions

∙ The number of electrons in a stable cation or anion = #electron in nearest noble gases

•Topic 7: Avogadro’s number

∙ Avogadro’s number: 1 mole = 6.022 x 1023 particles

∙ Ex: 1 mol carbon-12 = 6.022 x 1023 atoms of carbon-12

∙ Particles could mean atoms, molecules, ions, etc.  

∙ Moles to number of atoms

∙ to convert from moles of an atom to number of atoms, multiply the moles  by Avogadro’s number



∙ to convert from number of atoms to moles, divide the number of atoms by Avogadro’s number

∙ Mass to moles using molar mass

∙ molar mass describes the mass of a substance in one mole of that  substance

∙ calculated using Avogadro’s number

∙ to convert from mass to moles, divide the mass by the molar mass ∙ Mass to numbers of atoms

∙ mass to number of atoms: divide mass by molar mass to get moles, then  multiply by Avogadro’s number

∙ number of atoms to mass, divide number of atoms by Avogadro’s number  and then multiply by molar mass


•Topic 1: Types of bonding

∙ 2 main types

∙ Ionic bonding

∙ Occurs when a metal reacts with nonmetals to form ionic compounds ∙ Ionic bond = “an electrostatic attraction between an anion and a cation” ∙ An ionic compound is not a molecule, but a 3D lattice of alternating  cation and anion

∙ Covalent bonding

∙ occurs between two nonmetals or an nonmetal and metalloids ∙ pair of electrons shared between two atoms in a covalent bond •Topic 2: Types of formulas  

∙ Chemical formula – all compounds have a chemical formula; describes the  relative number of atoms/ions in the compounds

∙ Molecular formula – only for covalent compounds; describes the number of  atoms of each element in the compound

∙ Empirical formula – simples ratio of number of atoms

∙ For ionic compounds, the chemical formula = empirical formula •Topic 3: Types of models

∙ Ball and stick model

∙ Space filling model



•Topic 4: Naming ionic compounds and molecular  compounds

∙ Naming Ionic compounds  

Binary compounds

∙ Two elements

∙ Write out the cation and the anion with their charges

∙ “cross” the charge of the cation as a subscript on the anion and the  charge of the anion as a subscript on the cation

∙ Simply the ratio of cation to anion if applicable

∙ Ex: Ca2+ S2- -> Ca2S2 -> CaS

Compounds containing polyatomic ions

∙ similar process to naming binary ionic compound

∙ treat the polyatomic ion as an anion or cation

∙ if the polyatomic ion has a subscript, place parenthesis around it ∙ Ex: Ca2+ PO43- -> Ca3(PO4)2 

∙ Naming covalent compounds

∙ know the prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona,  deca

∙ the first element would not be called mono

 ex: CO2 is not monocarbon dioxide, it is just carbon dioxide

∙ name the first element with its prefix and name and the second  element with the prefix + element name – “ending” + “ide”

 ex: H2O becomes dihydrogen monoxide

•Topic 5: Formula/Molar Mass

∙ Formula mass = molecular mas = molecular weight = molar mass of a  compound

∙ Multiply the number of atoms by the molar mass of that atom and add all of  the molar masses of the various atoms together

•Topic 6: Mass Percent Composition

∙ Percent mass of element = [(mass of element * number of moles of elements in the compound)/ (mass of one mole of compound)] * 100%

∙ Mass percent can be used a conversion factor to go from mass of a  compound to mass of a particular element in a compound

•Topic 7: Molecular Formula vs Empirical Formula ∙ To determine the chemical composition of a compound using experimental  data

∙ NxOy decomposes into 24.5 g N & 70.0 g O

∙ Plan: mass N -> mol N  

 mass O -> mol O



then get the simple whole number ratio of x:y  

∙ The resulting formula is the empirical formula

∙ To get the chemical formula from the empirical formula, you need the  molar mass of the compound

∙ Divide the molar mass of the compound by the mass of the empirical  formula

∙ Then, multiply the resulting number with the empirical formula to get  the the chemical formula

•Topic 8: Balancing Chemical Equations

∙ use states (s) (l) (g) (aq)

∙ total number of atoms n in reactants = total number of atoms in products ∙ Combustion  

∙ Burning in O2(g)

∙ Other reactant is hydrocarbon, which contains C and H

∙ Start by balancing C

∙ Then balance H

∙ Balance O last

∙ Equations containing ionic compounds and polyatomic ions

∙ Treat polyatomic ions as 1 unit


•Topic 1: Stoichiometry, Mole-to-Mole and Mass-to-Mass  Conversions

∙ Reaction stoichiometry – numerical relationship between relative amount of  moles of all reactants and products in a balanced chemical equation ∙ Ex: combustion of C4H10(g)

∙ 2C4H10(g) + 13O2(g) -> 8CO2(g) + 10H2O(g)

∙ coefficients specify the relative amounts in moles of all reactants and  products

∙ mole ratio: 2 moles C4H10 : 13 moles O2 : 8 moles CO2 + 10 moles H2O ∙ These mole ratios can be used as conversion factor to perform calculations ∙ Mass-to-mass conversions:

∙ A + B -> C + D

∙ Mass A -> Moles A -> Moles C -> Mass C

∙ Use molar mass to convert between mass A and moles A

∙ Then, use mole ratio of A to B to convert from moles A to moles C ∙ Then, use molar mass to convert between moles C and mass C



•Topic 2: Limiting Reactants, Theoretical and Percent Yield ∙ The limiting reactant (LR) is the one that yields that the fewer moles of  product  

∙ To calculate which reactant is the limiting reactant,  

∙ Calculate how much product results from each of the reactant using  stoichiometry

∙ The lower yield is the actual theoretical yield

∙ The reactant that yields the fewer amount of product will be the  limiting reactant

∙ Percent yield  

∙ In most chemical reactions, the theoretical yield is not equal to the  actual yield

∙ Reasons

 Impurities

 Product may get washed away when purified

 Side reactions can occur

∙ If percent yield is above 100%, the product is wet or contains  impurities

∙ Percent yield = [actual (experimental) yield/theoretical yield] * 100%

•Topic 3: Solutions

∙ Substances are called AQUEOUS SOLUTIONS

∙ Amount of SOLUTE (measured in moles) is less than amount of  SOLVENT(measured in Liters)

∙ MOLARITY = concentration of solutions

 Molarity =(moles of solute)/(Liters of solution)

 Volume of solvent does not equal volume of solution

 Units =mol/L or M

 Ex: “the concentration of the solution is 2 M”

∙ Concentrated solutions contain more solute than dilute solutions  Concentrate solutions = those whose molarity is greater than 1  M

•Topic 4: Types of Aqueous Solutions and Solubility ∙ If a solid is an ionic compound, the attractive forces between the ions must  be broken in order to make a solution out of the solid

∙ A solid will dissolve when  

 Solute-solvent interactions > solute-solute interactions

 Solid forms complex ions with water molecules

∙ Conductivity

 When ions dissolve in water they move around  



 The movement of the charged particles cause ELECTRICAL  


 ex: “If NaCl conducts electricity, it is called electrolytic”

 Most ionic compounds that dissolve in water form ELECTROLYTES ∙ Covalent compounds in water

 No covalent bonds broken

 OH bonds can interact with water molecules

 No ions in solutions = no conductivity -> called NON-ELECTROLYTES ∙ 2 types of electrolytes

1. strong  

 good conductor

 high concentration of ions

 ex: strong acids - memorize - HCL, HBR, HI, HNO3, H2SO4, HClO4 2. weak

 poor conductor of electricity

 low concentration of ions

∙ Solubility of Ionic Compounds

o Not all ionic compounds are soluble in water, many are INSOLUBLE Ex: limestone

o Use the Solubility Rules (pg. 147 in textbook)

•Topic 5: Diluting a Solution

∙ Concentrated solution + more water = diluted solution

∙ STOCK SOLUTION = concentrated solution that is to be diluted before use ∙ DILUTION EQUATION: M1V1 =M2V2

where M1 is the initial molarity, V1 is the initial volume,  

 M2 is the final molarity and V2 is the final volume

∙ Solution stoichiometry

o A + B -> C +D  

o When you know molarity and volume of A or molarity and volume of B,  get the moles of A and B  

o Then, calculate moles of C or D like usual

•Topic 6: Precipitation Reactions

∙ Soluble salt (aq) + soluble salt (aq) -> precipitate (s) + soluble salt (aq) ∙ A PRECIPITATE forms upon mixing two solutions

∙ Know which is precipitate by looking at solubility rules

∙ Procedure for writing equations for precipitation reactions (molecular  equation)

1) Write the formulas of the solutions being mixed



2) Below the equation, write the formulas of the products that would  form

3) Look at solubility rules to see if any of the resultant products are  soluble

4) If all products are soluble, write no reaction

5) If there is an insoluble product, write the solid with (s) and any  soluble products with (aq)

6) Balance the equation

•Topic 7: Molecular, Ionic, and Complete Ionic Equations ∙ They are all ways of representing aqueous reactions

∙ Molecular equation - shows complete neutral formulas for each compound in  the reaction as if they existed as molecules

∙ Complete ionic equation - equation that individually list all ions present  o These better represent the dissociated nature of dissolved ionic  compounds

∙ Net ionic  

o SPECTATOR IONS appear unchanged on both sides of the equations  These are omitted from net ionic equation

 Only show the species that actually change during the equation

•Topic 8: Acid-Base Reactions


∙ GAS-EVOLUTION REACTION - a subtype of neutralization reaction in which gas forms

∙ Arrhenius definition: Acid produce H+ and Bases produce OH-  ∙ Strong acids completely dissociated in solution  

∙ Weak acid do not completely ionize in solution

∙ POLYPROTIC acid - contain 2 or more H+ ions

 Ex: H2SO4

∙ Strong bases completely ionize in solution - strong ELECTROLYTES ∙ Weak bases partially ionize in solution

 These are not written as ions when writing net ionic equations ∙ Acid+base -> water + a salt

∙ Acid-base TITRATION

 Analytical technique to determine the molarity of an acid which  is unknown

∙ Gas-forming reactions

 Acid + base -> gas + water + salt

 ex: 2HCl(aq) + NaCO3(aq) -> CO2(g)+ NaCl(aq)+ H2O(l)





∙ Air = homogeneous mixture of gases

∙ Atoms/ molecules are very fast, in random motion of varying speed,  separated by large distances and compressible  

•Topic 1: Gas Pressure

∙ Pressure = Force/ area  

∙ Atmospheric pressure = 14.7 lbs/ in2 = 14.7 psi = 110,325 N/m2 = 1 atm ∙ Conversion units

 1 atm = 760 mm Hg

 1 mm Hg = 1 torr

 1 atm = 101, 325 Pa

 1 atm = 14.7 psi

∙ As amount of gas increases, pressure increase

∙ As amount of gas in a vessel decreases, pressure decreases ∙ Measuring gas pressure

 Use a barometer

 Torricelli discovers mercury barometer

o Height of column Hg = atmospheric pressure  

measured in mm

 Manometers, LAB quest system

•Topic 2: Gas Laws

∙ 4 physical properties related to pressure - Pressure, volume, temperature,  amount  

∙ 1) BOYLE’s LAW

• P and V are inversely proportional  

• If P increases, V decreases and if P decreases, V increases and vice versa • Only holds true if T and n are constant  


• P1V1 = P2V2

∙ 2) Charles’ Law

• T and V are directly proportional

• If T increases, V increases and if T decreases, V decreases and vice versa • Only holds true if P and n are held constant

• Temperature must be in Kelvins!

•K = ℃ + 273

• V1/P1 = V2/T2

∙ 3) Avogadro’s Law

• V and n are directly proportional

• If n increases, V increases and if n decreases, V decreases and vice versa • Only holds true if P and T are held constant

• V1/n1 = V2/n2

∙ 4)Amonton’s Law

• P and T are directly proportional  

• If P increases, T increases and if P decreases, T decreases and vice versa • Only holds true if V and n are held constant

• P1/T1 = P2/T2

∙ 5) The ideal gas law

• Constant R = .08206 L atm/mol K  

• PV = nRT 

• Make sure P is in atm, V in L, n in moles, T in K

•Topic 3: Molar Volume and Density of Gases ∙ Definition The volume of one mole of nay gas

∙ Changes depending on P and T  

∙ Standard Temperature and Pressure = STP

 T = 273 K

 P = 1.00 atm

∙ The volume of 1 mol of gas at STP = 22.4 L

∙ Density of gases is measured in g/L

∙ Density is proportional to molar mass

∙ PV = nRT

 PV = (mRT)/(molar mass)

 P(MM) = (m/v) RT


•Topic 4: Mixtures of Gases and Partial Pressure ∙ Dalton’s Law of Partial Pressures

• Ptot= P1 + P2 + P3 +… 

∙ To calculate partial pressure:

• Need mole fraction, X

• X = moles of gas A/ total moles of gas mixture

• Ex: XA = nA/(nA + nB)

• Partial pressure of gas A, PA = XA * Ptot

•Topic 5: Collecting Gases over Water

∙ Often, when the product of a chemical reaction is a gas, it is collected by  displacement of water

∙ As the gas forms, it bubbles through water and gathers in the collection false ∙ The gas collected in this way is not pure because it is mixed with water vapor  ∙ The partial pressure of the water in the resulting mixture is its vapor  pressure, which depends on the temperature of the water  

∙ Therefore, Pgas = Ptotal - Pwater

•Topic 6: Gases and Stoichiometry

∙ Plan: mass A -> moles A -> moles B -> mass B

∙ Plan for gases: P, V, T of gas A -> moles gas A -> moles gas B -> P, V, T of  gas B

•Topic 7: Kinetic Molecular Theory

∙ Definition: a model that explains the properties and behavior of gases ∙ Developed after the gas law, explains them

∙ A gas is modeled as a collected of particles in constant motion; a single  particle movies in a straight line until it collides with another particles ∙ The three postulates (assumptions)

1) the size of a particle is negligibly small

 gas particles cannot occupy any volume  

 this assumption is justified because under normal pressures, the  space between particles in a gas is very large compared to the size of the particles themselves

2) average kinetic energy of a particle is proportional to the temperature in  kelvins

 motion of gas particles is due to thermal energy


 the higher the temperature, the faster the overall motion and the  greater the average kinetic energy

 kinetic energy (1/2 * mv2)

 note that kinetic energy, NOT volume is proportional to  


 two different samples can have same kinetic energy and very  different velocities

3) collision of one particle with another or with the walls is completely elastic  elastic = no overall loss energy

 particles have no attraction forces between each other

∙ Relation to the gas laws

1) Boyle’s Law

 if V decreases, number of collisions increase

 if number of collisions increase, pressure increases

2) Charles’ Law

 If T increases, KE increases, no of collisions increases  

 If P is constant, gas will have to expand and V will have to increase 3) Avogadro’s Law

 If n increases, number of collisions increases

 P is constant, therefore V must increase

∙ Temperature and Molecular Velocity

o KE = 1/2 * mv2 

o Lighter particles travel faster than heavier particles

o For velocity, use ROOT MEAN SQUARE VELOCITY (urms)

 urms = √U2 

 We use this formula because of the wide range of velocities in a  sample of gas as it gives a more accurate average

 urms = √(3RT)/(MM) 

 Use molar mass in kg/mol !

 R = ideal gas constant = 8.314 J/mol K

 1 J = 1 kg m^2/s^2

o As T increase, root mean square velocity increases and as T decreases, root mean square velocity decreases

o At constant temperature, as MM increases, root mean square velocity  decreases and as MM decreases, root mean square velocity increases

•Topic 8: Diffusion and Effusion

∙ Molecules collide with each other in elastic collision but change directions after they collide

∙ Average distance between collision = MEAN FREE PATH

∙ As P increases, collisions increase and mean free path decreases  17

∙ Diffusion

o The process by which gas molecules spread out

o Light molecules travel faster than heavier molecules

o Mean free path of lighter molecules is greater than mean free path of  heavier molecules

∙ Effusion

o The process by which gas molecules escape from a container through a  small hole and into a vacuum

o 2 different gases, A and B

o GRAHAM’s LAW OF EFFUSION Rate a/ rate b = √(MM b/ MM a)

•Topic 9: Real Gases

∙ There is no such thing as an ideal gas

∙ For all gases, the 3 postulates of the kinetic molecular theory do not hold true o why?

o At high P, size of gas particles are significant  

o Collisions b/w gas particles are non-elastic because attractive and  repulsive forces are present  

∙ Behavior of gases at extreme conditions

o At high pressures and low temperatures, we cannot use the ideal gas  equation

o High pressure - volume of gas particles is significant and they take up  some room

o Low temperature - temperature decrease, root mean square velocity  decreases, fewer collisions and collisions occur with less kinetic energy • therefore, molecules may interact with each other and attraction  forces occur between them and the attractive force affect the  


• The result is a decrease in the number of collisions which leads to a  decrease in the pressure compared to that of an ideal gas  



∙ themo=heat

∙ Heat is a from of energy

∙ Most chemical reactions involve heat

∙ Heat can be absorbed - this type of reaction would be called ENDOTHERMIC or  heat can be released - this type of reaction would be EXOTHERMIC ∙ Thermochemistry - the study of the relationships between chemistry and energy

•Topic 1: Energy - Key Definitions

∙ ENERGY - the capacity/ ability to do work

∙ WORK is the result of a force acting through a distance

∙ KINETIC ENERGY is energy due to motion

∙ Heat energy = THERMAL ENERGY

∙ POTENTIAL ENERGY = stored energy  

• In chemistry, potential energy is chemical energy  

• i.e., the energy associated with the chemical bonds as well as the  relative positions of electrons, protons, neutrons

• Chemical energy is stored in the bonds

∙ Breaking bonds is endothermic - requires the absorption of energy  ∙ Making bonds releases energy and is exothermic  

∙ LAW OF CONSERVATION OF ENERGY - FIRST LAW OF THERMODYNAMICS • Energy cannot be created or destroyed

• Can be transferred from one body to another


• System - anything that is under investigation

• Surroundings - everything with which the system can exchange energy ∙ Units of energy

• Joule (J) = 1 kg m2/s2 

• calorie (cal) = 4.184 J

• Calorie (Cal) = 1 kcal = 1000 cal

• Kilowatt-hour (kWh) = 3.60 X 106 J

•Topic 2: The First Law of Thermodynamics  

∙ The total energy of the universe is constant  



o E = kinetic energy + potential energy of all particles in a system o Ex: in a water molecule, there is translational energy, vibrational  energy, rotational energy and energy in the bonds

o Internal energy is a STATE FUNCTION, I.e., value depends only on the  present state of the system, now how the system arrived at the state  ex: mountain climber analogy - no matter which path an climber  takes, it only matters that the altitude will be the same at the  


∙ Changes in internal energy, ΔE

o ΔE = Efinal - Einitial = Eproducts - Ereactants

∙ 1st Law: ΔEsurroundings = ΔEsystem

∙ ΔEtotal = 0 = ΔEsurroundings + ΔEsystem

∙ If the reactants have a higher internal energy than the products, ΔEsys is  negative and energy flows out of the system into the surroundings ∙ If the reactants have a lower internal energy than the products, ΔEsys is  positive and energy flows into the system from the surroundings

•Topic 3: Heat and Work

∙ When energy is transferred from system to surroundings or vice versa, heat  energy and work is done

∙ According to first law of thermodynamics, the change in the internal energy  of the system must be the sum of the heat transferred (q) and the work done  (w)

o ΔE = q + w 

o q…can be + or ­

o w…can be + or ­

o Signs conventions:

 When +: q - system gains energy from the surroundings

o w - work is done on the system by the surroundings

o ΔE ­ energy flow into the system , system gains E, endothermic

 When -: q - system loses energy to the surroundings  

o w - work is done by the system on the surroundings

o ΔE ­ energy flows out of the system, system loses E, exothermic 

∙ HEAT is the transfer of thermal energy from system to surroundings o Temperature is a measure of the thermal energy of either the system  or surroundings

o When a system loses or gains heat, there is always a change in  temperature = ΔT

o Q = C ΔT 

 C is a constant that describes the heat capacity of a system,  units: J/ ℃ 

 more commonly, Specific heat capacity, Cs, is used 


 SPECIFIC HEAT CAPACITY is the amount of heat require to raise  the heat of 1 g of the substance by 1℃, units: J/ g ℃ or J/ g K 

 Q = mCsΔT  

∙ WORK is done by the system on the surroundings or by the surroundings on  the system  

o Type of work -> PRESSURE-VOLUME WORK = expansion work  o W = -PΔV 

o Work is usually for the system (don’t really worry about the surroundings) 

•Topic 4: Measuring ΔE  

∙ Measuring ΔE at constant volume 

o ΔE = q + w, but at constant V, w = 0, therefore at constant V, ΔE = qv (subscript v denotes at constant volume)

o Method: bomb calorimeter  

 Sample(usually food) is combusted -> heat water -> cause ΔT  qcal = Ccal ΔT 

∙ Ccal = heat capacity of calorimeter (J/ ℃ or kJ/ ℃)

 Heat lost by food combusting = heat gained by a calorimeter  qrxn = -qcal

 qrxn = -Ccal ΔT 

∙ Measuring ΔE at constant pressure 

o Most chemical reactions are carried out at constant P rather than  constant volume 

o ENTHALPY = ΔE at constant pressure, denoted by H 

o H = E + PV 

o Changes in enthalpy, ΔH = ΔE + ΔPV 

 At constant P, ΔH = ΔE + P ΔV 

 ΔH = qp (subscript p denotes at constant pressure) 

o Enthalpy is the heat absorbed or released in a chemical reaction at  constant P 

 Endothermic - when ΔH is positive and heat is absorbed 

 Exothermic - when ΔH is negative and heat is released 

o Method: coffee cup calorimeter  

 ΔT = ΔHrxn

 Heat lost or gained by a solution, qsol’n

 qsol’n = mass sol’n * Cs sol’n * ΔTsol’n  

 Heat lost by rxn = heat gained by surroundings (sol’n) 

∙ +qrxn = -qsurr

 Difference between qrxn and qsurr is that qrxn doesn’t account for  spectator ions and water 

 Calc qrxn in J -> kJ -> kJ/mol -> ΔHrxn


∙ ΔHrxn and stoichiometry 

o Every rxn has its own value go ΔHrxn

o ΔHrxn can be used as a conversion factor  

o ex: -1648 kJ/ 4 mol Fe 

o when calculating mass, make sure to leave out the negative sign in  heat  

• Topic 5: Hess’s Law and Other Relationships Involving  ∆Hrxn

Lecture 20 and 21 (10/18/17 and 10/20/17), Section 6.8 in Tro

∙ 3 quantitative relationships between a chemical reaction and ∆Hrxn  

o ∆Hrxn is also multiplied by some factor if a chemical reaction is  multiplied by that factor

o ∆Hrxn changes sign when the chemical reaction is reversed

o When a chemical reaction can be expressed as a sum of a series of  steps, ∆Hrxn for a chemical reaction is the overall sum of the heats of  reactions for each step of the series – this principle is known as Hess’s  Law 

∙ Use of Hess’s Law: Determine ∆H for a reaction without having to directly measure it in  the lab

∙ Tips for doing a problem with Hess’s Law

o We’ll call the final equation we’re trying to achieve “the equation of interest”

o If the equation of interest has a reactant that is listed as a product in  one step, reverse the step  

o If the equation of interest requires two moles of a reactant and a step  has only one mole of that reactant, multiply the step by two


o After you have manipulated all the equation as necessary, calculate  the ∆Hrxn of each of the steps using the quantitative relationships listed  above

o Then add all the ∆Hrxn of each step to get the final ∆Hrxn of the equation  of interest

Topic 6: Enthalpies of Reaction from Standard Heats  of Formation

Lecture 21 and 22 (10/20/17 and 10/22/17), Section 6.9 in Tro ∙ ∆Hrxn can also be calculated using standard enthalpies of formation ∙ ∆Hrxn is the change in enthalpy for a chemical reaction  

o i.e. the difference in enthalpy between the products and the reactants  

∙ ∆Hrxn is a standard state function, so its value depends on the initial and final  values, not the pathway taken  

o therefore, we can relatively define the zero of enthalpy conveniently  ∙ Standard for enthalpy  

1) Standard state

∙ For a gas, standard state is the pure gas at 1 atm

∙ For a liquid or solid, standard state is the pure substance in its most stable form at 1 atm and temperature of interest (25 degrees  

Celsius usually)

∙ For a substance in solution, the standard state is concentration of 1  M

2) Standard Enthalpy Change (∆H0)  

∙ The standard enthalpy change is the change in enthalpy for a  process when all reactants and products are in their standard state


∙ ∆H0rxn = ∆H0products - ∆H0reactants  

3) Standard Enthalpy of Formation (∆H0f)

∙ Every compound and element has its own ∆H0f – these are given in  table 6.2 of Tro 

∙ For a pure compound, the standard enthalpy of formation is the  change in enthalpy when 1 mole of the compound forms from its  constituent elements in their standard states

∙ For a pure element in its standard state: ∆H0f = 0

∙ For an element not in its most stable form, ∆H0f ≠ 0

o Eg: carbon in diamond form

∙ Calculating Standard Enthalpy Change for a Reaction

o Consider these principles:

 Formation of a compound from its constituent elements in their  standard states = standard heat of formation (∆H0f)

 Decomposition of a compound into its constituent elements =  negative of the standard heat of formation (∆H0f)

o Now:

1) Decompose the reactants into their constituent elements in their standard states

Reactants -> Elements ∆H1 = - ∑∆H0f (reactants)

2) Form the products from the constituent elements in their  

constituent elements

Elements -> Products ∆H2 = + ∑∆H0f (products)

∑ = “the sum of”

∆H1= sum of the negatives of the heats of formation of the  


∆H2=sum of the heats of formation of the products

Reactants -> Elements ∆H1 = - ∑∆H0f (reactants)

 +Elements -> Products ∆H2 = + ∑∆H0f (products)


Reactants -> Products ∆H0rxn = ∑∆H1 - ∑∆H2

∙ Overall, to calculate ∆H0rxn , subtract the heats of formation of the reactants  multiplied by their stoichiometric coefficients from the heats of formation of the  products multiplied by the stoichiometric coefficients

∙ ∆H0rxn = ∑np∆H0 f (products)- ∑np∆H0 f (reactants) 



• Topic 1: The Nature of Light (Lecture 22 – 10/25/17;  7.2 in Tro textbook)

Wave-particle duality of light – certain properties are best described by  thinking lights as a wave, other properties are best described by thinking  of it as a particle

Wave Nature of Light

o Light is not matter – has no mass and no volume

o It is ELECTROMAGNETIC RADIATION, a type of energy with  

oscillating electric and magnetic fields

o SPEED OF LIGHT (c) = 3.00 X 108 m/s

o Electromagnetic waves have amplitude and wavelength  

o Amplitude – vertical height of a crest or depth of a trough and  determines the intensity or brightness of the light;

 Greater the amplitude, greater the brightness of light

o Wavelength (λ) – distance in space between adjacent crests or troughs;  measured in units of distance; wavelength of a light determines its  color

o Frequency (√) – number of cycles that pass through a stationary point in a given period of time, units = s-1 

o Frequency is inversely proportional to the wavelength  

 ν λ  = c/ 

The Electromagnetic Spectrum

o Visible light is one part – many other types of EM radiation

o Figure 7.5 – memorize in order of lowest energy to highest energy: 

Radio, microwave, infrared, visible, ultraviolet, x-ray,  


Visible: Red, orange, yellow, green, blue, indigo, violet

o Interference and Diffraction (pg. 268 – 269)

 Waves interact in a characteristic way called interference

 Can either cancel each other out or build each other up

 Constructive interference: waves of equal amplitude from  

two sources align with overlapping crests to products a wave  

with twice the amplitude

 Destructive interference – waves are out of phase and align so 

that the crest from one source overlaps the trough from the  


other source and the waves cancel by destructive  


 Diffraction – when a wave encounters an obstacle or a slit that  is comparable in size to its wavelength, it bends around it 

 Each slit acts as a new wave source, and the two new waves interfere with each other 

 Pattern that results – has a series of bright and dark lines that can  be viewed on a screen  

 Waves out of phase (destructive) make a dark spot 

 Waves in phase make bright spot  

 “when a beam of light passes through two small slits, the two  resulting waves interfere with each other. Whether the  

interference is constructive or destructive at any given point  

depends on the difference in the path lengths traveled by the 

waves. The resulting interference pattern can be viewed as a  

series of bright and dark lines on film”  

Particle Nature of Light

o Key discoveries brought the classical wave nature of light into  question

o Photoelectric effect – light of a minimum energy will cause the  ejection of electrons

 Study particle nature for light

 Amount of energy must exceed the electron’s binding energy

 Light used to dislodge electrons exhibits a threshold  

frequency, below which no electrons are emitted from the  


 Einstein’s theory – light is lumpy 

 light comes in pockets called photons or quantum

o The amount of energy in photon or a quantum of light

 E=h √ 

 h = Planck’s constant = 6.626 X 10-34 J s  

 E = (hc)/ λ 


• Topic 2: Atomic Spectroscopy and the Bohr Model  (Lecture 22 and 23 – 10/25/17 to 10/27/17, 7.3 in Tro) ∙ Atomic spectroscopy, the study of electromagnetic radiation absorbed  and emitted by atoms, was important in suggesting a wave nature for  particles  

∙ An atom reemits energy as light when an atom absorbs energy  ∙ Emission spectrum – a series of bright lines that occurs when light emitted by  an element is separated by passing it into a prism

 Emission spectrum of a certain element is always the same


 Therefore, analysis of light allows us to identify elements

 As number of electrons in an element increase, the number of  lines in an emission spectrum increases

∙ Rydberg developed an equation to calculate wavelength of each line in  the hydrogen emission spectrum

 1/ λ = R (1/m2 – 1/n2)  

 m and n are whole numbers

 m < n

 R = Rydberg constant = 1.0997 X 107 m-1 

 Neils Bohr explained the reason behind this equation  

∙ Bohr model of the atom

 Electrons travel in orbits around the nucleus

 Energy of each orbit is fixed and can only have certain values   Electrons can be excited into higher orbits when they absorb  energy (heat or electrical)

 When the electron jumps back down, energy is released in the  form of IR, visible or UV light

 Bohr model only explains behavior of electron in H atom  

• Topic 3: The deBroglie Principle (Lecture 23 –  10/27/17, 7.4 in Tro)

∙ Electrons behave like light waves  

∙ Undergo diffraction

∙ E=mc2 and E=hc/ λ 

∙ Therefore, λ = h/mv

∙ h in J *s, m in kg, v in m/s

∙ deBroglie’s equation only works in the atomic world when particle have very  small mass and travel at high speeds

• Topic 4: Heisenberg’s Uncertainty Principle and  Indeterminacy (Lecture 24 – 10/30/17, 7.4 in Tro) ∙ You cannot determine simultaneously the wave nature and particle  nature of the electron

∙ Therefore, you cannot know the exact position and exact velocity  simultaneously

∙ Heisenberg’s Uncertainty principle 

 ∆x * m∆v > h/(2π)

 ∆x is uncertainty in position  

 ∆v is uncertainty in velocity

∙ Uncertainty led chemists to think about probability  

 The probable location of the electron in the atom is where it is  MOST LIKELY to be found

 Probability distribution map – a statistical map that shows where an  electron is likely to be found under a given set of conditions  


∙ Indeterminacy – the future path of an electron is indeterminate and can only be described statistically

• Topic 5: Quantum Mechanics and the Atom (Lecture  24 – 10/30/17, 7.5 in Tro)

∙ Orbital – a probability distribution map showing where the electron is  likely to be found  

∙ Schrodinger took the ideas of Planck, Einstein, deBroglie, Heisenberg  and developed an equation for the energy of an electron in an atom

Hψ= Eψ;

∙ H is the Hamiltonian operator – represents the total energy of the electron in  the atom

∙ E is the actual energy of the atom  

∙ Ψ is the wave function, a mathematical function that describes the  wavelike nature of electrons  

∙ Ψ2 represents an orbital, a position probability distribution of the  electron  

∙ Quantum Numbers  

o Electrons can be present in different orbitals  

o Several types of orbitals, different energies – they depend of  


o Quantum numbers are what specify each orbital 

1) The Principle Quantum Number, n  

∙ Always an integer

∙ determines the overall size and energy of an orbital

∙ Possible values: n=1,2,3….

∙ For a H atom, the energy of an electron in an orbital with  

quantum number n is  

En= -2.18 X 10-18 J (1/n2)  

∙ The energy of the electron when it is very far away from the  atom is taken to be zero

∙ Therefore, the energy is negative because the energy of the  electron in the atom is less than the energy of the electron when it is very far from the atom

∙ The more negative the energy, the more stable the electron is  ∙ As n increases, the spacing between energy levels decreases  2) The Angular Momentum Quantum Number, l  

∙ Always an integer

∙ Determines the shape of the orbital  

∙ Possible values: 0, 1, 2,…(n­1)

∙ Ex; when n = 1, l cannot equal 1


∙ Values of l are assigned letters

o l = 0, s

o l = 1, p

o l = 2, d

o l = 3, f

∙ letters correlate to line spectra of each given orbital  

o s has a sharp line spectra  

o p has a principled line spectra

o d has a diffused line spectra, and so on

3) The Magnetic Quantum Number, ml

∙ Always an integer

∙ Specifies the orientation of the orbital with respect to x, y and z  axes

∙ Possible values: range from –l to +l  

∙ Ex: if l = 1, the possible values of ml are -1, 0, 1

∙ One only type of s orbital

∙ Three types of p orbital: px, py, pz

∙ Five types of d and seven types of f

4) The Spin Quantum Number (mx)

∙ Specifies the orientation of the spin of the electron  

∙ Electron spin is a fundamental property of an electron

∙ All electrons have the same amount of spin  

∙ Orientation of the spin is quantized

∙ Two possibilities only: spin up (+1/2) or spin down (-1/2)

• Topic 6: Atomic Spectroscopy Revisited (Lecture 25 –  11/1/17, 7.5 in Tro)

∙ Each wavelength in the emission spectrum of an atom corresponds to  an electron transition between quantum-mechanical orbitals  

∙ When an electron absorbs energy, it is excited or promoted from a  lower-energy level orbital to a higher-energy-level orbital  

∙ When the electron is excited, however, it is unstable

∙ Therefore, the electron quickly falls back to a lower energy orbital  ∙ When it falls back, it releases a photon of light containing an amount of energy precisely equal to the energy difference between two energy  levels  

o The difference in energy between two levels ninitial and nfinal is  given by ∆E = Efinal – Einitial  

∙ Depending on the energy level it falls to, the electron emits either UV,  visible light or IR  

∙ Formula:  

∆E = Efinal – Einitial


∙ Law of conservation of energy

∆Eatom = - ∆Ephoton

∙ The transitions between orbitals that are farther apart in energy  produce light that is higher in energy than transitions between orbitals  that are closer together

• Topic 7: Shapes of Orbitals (Lecture 24 – 10/3/17, 7.6  in Tro)

∙ Shapes of atomic orbitals are important because covalent chemical  bonds depend on the sharing of the electrons that occupy these orbitals ∙ Ψ2= probability density

∙ Orbital is a mathematical representation of where the electron is most  likely to be, 90 % of the time

∙ S-orbital (l = 0)

o Spherically symmetrical orbital  

o To get a better idea of where the electron is most likely to be found, a plot called the radial distribution function which shows the total  probability of finding the electron with a thin spherical shell at a  distance r from the nucleus  

o At the nucleus, r = 0, total radial probability is zero  

o A node is a point in which the wave function (Ψ), probability density, and radial distribution function all go to zero

 2s has 1

 3s has 2, and so on  

o Shape primarily determined by l

∙ P-orbital (l = 1)

o Not spherically symmetric  

o Instead have two lobes of electron density on either side of the  nucleus and a lobe at the nucleus

o ml = -1, 0, +1  

 therefore, 3 different p orbitals  

∙ d-orbital (l = 2)

o ml = +2, +1, 0, -1, - 2

 therefore, five different d orbitals  

 four have a cloverleaf shape with four lobes of electron  

density and the last is a donut-shaped ring  

∙ f-orbital (l = 3)  

o ml = -3, -2, -1, 0, +1, +2, +3

o more lobes and more nodes than d-orbital  



• Topic 1: Electron Configuration (Lecture 25 – 11/1/17,  8.2 in Tro)

∙ Arrangement of the electrons in an atom

∙ Return to quantum numbers

 4th quantum number, ms – spin quantum number, describes the spin  of the electron

 only two possibilities for ms - +1/2, -1/2  

 +1/2 is clockwise spin (up in orbital diagram); -1/2 is  

counterclockwise spin (down in orbital diagram)

∙ electrons have opposite spins when in the same orbitals

 Pauli exclusion principle – no two electrons have the same set of  four quantum numbers

• Topic 2: Sublevel splitting

∙ In general, Es orbital < Ef orbital < Ed orbital < Ef orbital

∙ But for He and all other elements, there is energy level splitting ∙ (degenerate = orbitals of same energy)

• Topic 3: Electron Configuration and the Multielectron  Atom

∙ Aufbau principle – orbitals of lower energy are filled before orbitals of higher 


 Think building up

∙ Know electron configuration up to calcium

∙ Noble gas/short hand notation

 Ex: Na 1s2 2s2 2p6 3s1 

 Na [Ne] 3s1 

∙ Hund’s rule – when filling orbitals of the same energy electrons fill them singly at

first, with parallel spin

 I.e., electrons prefer to be unpaired if possible  


∙ Half full and full shells are extra stable  

∙ E3d < E4s

 Therefore, 4s electrons are lost first before 3d electrons when  atoms are ionized  

∙ D-block exceptions

 Copper, Chromium  

∙ F-block

 Really hard to predict

 For expected configuration, just count across  

 Even though the true configuration differs from these most of  the time

• Topic 4: Electron Configuration of Ions

∙ Chemical properties of elements are determined by the number of  valence electrons

∙ Elements can gain e-‘s and lose e-‘s and will tend to have the same number  of electrons as the nearest noble gas 

∙ Isoelectric – having the same number of electrons  

∙ Octet rule – having a full eight valence electrons is considered most stable ∙ Duet rule – for hydrogen, two valence electrons is stable

• Topic 5: Atomic Trends

∙ Atomic radii

 Down a group, radii increases – as n increase, size of orbitals 

increases, therefore r increases

 Across a period, radii decreases because nuclear charge  

increases and the nucleus pulls in the orbitals closer, therefore  radii decrease

 Zef (effective nuclear charge)  

∙ Ionic radii

 Electron configuration of the transition elements and their ions  After the 4s orbital is filled, the energy of the 3d orbital is below  

that of 4s

 Therefore, electrons in the 4s are removed first


 Transitions metals have magnetic properties due to unpaired  electrons

∙ Diamagnetic – all electrons paired and not attracted to outside 

magnetic force

∙ Paramagnetic – unpaired electrons and attracted to outside 

magnetic force

 Cations are smaller than their corresponding atoms

∙ With loss of electrons from orbitals of highest quantum  

number, the radius decreases

 Anions are larger than their corresponding atoms

∙ More electrons in an orbital = more repulsion between  

electrons = orbital increases in size

 Isoelectric ions – ions of different elements that have the same number  of electrons  

∙ Ionization energy, IE

 Ionization involves removing an electron completely from a  

gaseous atom  

 All ionization energies are endothermic  

 As atomic radii increases, ionization energy decreases because it

is easier to remove the outer electrons from larger atoms  

 Down a group, IE decreases  

 Across a period, IE increases but there are exceptions  

∙ i.e., hard to remove electrons from a stable, half-filled  


 trend in successive IE’s

∙ IE1 < IE2 < IE3 < IE4

∙ Electron affinity, En  

 Involves an atom gaining an electron  

 No regular trend in periodic table

 Group 17 (halogens) have the greatest (most negative) values of electron affinity

∙ Metallic characters

 Across a period, metallic character decreases

 Down a group, metallic character increases  



• Topic 1: Types of Chemical Bonds  

Lecture 28, 11-8-17; Section 2

∙ Chemical bonds form because they lower the potential energy between charged particles that compose atoms

∙ I.e., PE of atoms > PE of molecules or ions

1) Ionic

∙ Electrostatic attraction between metals and nonmetals

∙ Metal and nonmetal react to form a cation (metal loses e-) and  anion (nonmetal gains e-)

∙ Electrons transferred from metal (cation) to nonmetal (anion) ∙ When bonds forms, the ions are attracted to one another, lowering  their overall potential energy  

∙ For ionic compound, PE is directly proportional to q+ and q 

2) Covalent

∙ Forms between nonmetals

∙ Electrons shared between atoms

∙ When the bond forms, shared electrons interact with the nuclei of  both of the bonding atoms, lowering their potential energy  

∙ The molecule is stable when the repulsive forces are balanced by  the attractive forces

3) Metallic

∙ Occurs in metals

∙ Sea of electrons surrounds cations  

• Topic 2: Simple Lewis Dot Structures of Atoms

 (9.3, Lecture 29, 11-10-17)

∙ Represent electrons using dots; helps explain properties of elements ∙ Draw each valence electron as one dot

∙ Why valence electron? They are involved in chemical bonding  ∙ Lewis dot structures are a way to visualize the number of valence  electrons  

∙ Cannot draw Lewis dot structures of d-block or f-block elements 35

• Topic 3: Ionic Bonding: Lewis Symbols and Lattice  Energies  

 (9.4, Lecture 29, 11-10-17)

∙ Lewis model can also applied to ionic bonding

∙ Ionic bonding is represented by moving electron dots from Lewis  symbol of the metal to the Lewis symbol of the nonmetal.  

∙ Example:  

+ - 

Na Cl Na Cl

∙ Lattice Energy

o Energy associated with the formation of a crystalline lattice of  alternating cations and anions from the gaseous ions

o Lattice Energy is always negative because formation of the lattice is exothermic

o The more negative the lattice energy, the stronger the ionic bond ∙ Trends in Lattice Energy

o Ion Size – As radius decreases, the lattice energy becomes more  negative because the potential energy of oppositely charged ions  becomes more negative as the distance between ions decreases.  

o Ion Charge – As the product of the charges increases, the potential  energy increases. Therefore, lattice energy increases

∙ Effect of lattice energy on other properties

o Melting points – As lattice energy increases (becomes more  negative), melting point increases.

o Solubility – As lattice energy increases (becomes more negative),  solubility decreases.

• Topic 4: Covalent Bonding: Lewis Structures  (9.5, Lecture 29, 11-10-17)  


∙ Octet rule – When the atom has 8 valence electrons, it is fully stable. ∙ Duet rule – Exceptions for several atoms like hydrogen and helium;  when they have 2 valence electrons, they are stable.

∙ Bonding pair – A pair of electrons shared between two atoms. 

∙ Lone pair – A pair that is associated with only one atom. 

∙ Bonding pair are represented by dashes.  

∙ Double and Triple Covalent Bonds

o Two atoms may share more than one electron pair to get octets.  o Double bond – two electron pairs are shared between two atoms.  o Triple bond – Three electron pairs are shared between two atoms. 

• Topic 5: Electronegativity and Bond Polarity  (9.6, Lecture 29 – 11­10­17)

∙ When electrons are shared equally, the polarity has to be represented  in the L Lewis structure

 Either an arrow with a positive sign on the tail  

 Or delta plus and delta minus sign  

 The delta minus sign is put on the element that attracts the  electron.  

δ- δ+ 


∙ Electronegativity – the ability of atoms to attract electrons onto itself when part  of a molecule

 All electronegativity values are relative to that of fluorine

 Fluorine’s electronegativity value is 4.0

 Trends in electronegativity

 Decreases down a group

 Increases across a period  

∙ Polarity

 Difference in electronegativity determines if a bond between 2  atoms is non-polar, polar, or ionic  

 – 0.4 difference = nonpolar covalent

 0.4-2.0 difference = polar covalent

 2.0 – 3.3 difference = ionic

 when the difference is 2.0, you have to know the compound  


• Topic 6: Lewis Dot Structures of Polyatomic Ions and  Ionic Compounds

 (9.7, Lecture 30, 11-13-17)

∙ Steps

 Write the correct skeletal structure for the molecule. Remember, that  hydrogen atoms are always terminal atoms (not central). Also, place  the more electronegative elements in terminal positions and the less  electronegative in terminal positions.  

 Calculate the total number of electrons for the Lewis structure by  summing the valence electrons of each atom in the molecule.  

Remember, for polyatomic ions, consider the charge of the ion when  calculating the number of electrons.

 Distribute the electrons among the atoms, giving octets to as many  atoms as possible.

 If any atoms lack an octet, form double and triple bonds as necessary  to give them octets.  

• Topic 7: Resonance and Formal Charges

 (9.8, Lecture 305, 11-13-17)


∙ Resonance

 For some molecules, we can write more than one valid Lewis structure  The structures are called resonance structures

 The actual structure of the molecule is intermediate (an average)  between the two resonance structures, not a mixture of 50% of one and  50% the others. This structure is called the resonance hybrid

∙ Formal Charge

 The charge of an atom if all bonding electrons were shared equally  between the bonded atoms

 Not a real charge

 FC = # valence electrons – (# lone pair electrons + ½ # bonding  electrons)

 Sum of formal charges of all the atoms = overall charge of molecule  Small or zero formal charges on individual atoms are better (more stable)  When formal charge cannot be avoided, the negative formal charge  should be on the most electronegative atom

• Topic 8: Exceptions to the Octet Rule

 (9.9, Lecture 31, 11-17-17)

∙ Odd electron species – called radicals

 Total number of electrons is an odd number  

 One unpaired electron on atom

 Reactive and toxic

∙ Incomplete octets (less than 8 valence electrons on atom)

 Ex: BF3

 Could give fluorine a positive formal charge to give boron an octet, but  this is extremely unstable

• Expanded octet (more than 8 valence electrons on atom)

 Common among 3 period elements because the d-block is “energetically  accessible

 I.e. they are not much higher in energy so they can accommodate extra  electrons

 Look at formal charges to see which is the most favorable resonance  structure

• Topic 9: Bond Energies and Bond Lengths

 (9.10, Lecture 33, 11-25-17)


∙ Use of bond energies – individual bond energies can be used to estimate  enthalpy changes in reaction (∆H0rxn) when standard enthalpies of  formation for all the reactants and products of a reaction are not available  ∆H0rxn = ∑(∆H’s bonds broken) - ∑(∆H’s bonds formed)

 Forming bonds is exothermic, i.e. releases energy

 Breaking bonds is endothermic, i.e. requires energy

 The reaction is exothermic when weak bonds break and strong bonds  form (when ∆H0rxn is negative)

 The reaction is endothermic when strong bonds break and weak bonds  form (when ∆H0rxn is positive)

∙ Bond energy definition – The bond energy of a chemical bond is the energy required to break 1 mole of the bond in the gas phase.  

 Bond energies are always positive because it always takes energy to  break a bond.  

 The higher the bond energy, the stronger the bond.  

 Compounds with stronger bond energies tend to be more chemically  stable than compounds with weaker bond energies.  

∙ Average bond energy – an average of the bond energies for that bond in a  large number of compounds

∙ Factors which affect bond energy

 Type of atoms  

 Single vs multiple bonds

 Bond energy of triple bonds > bond energy of double bonds > bond  energy of single bonds for a particular pair of atoms  

∙ Bond lengths

 Bond length is the average length of a bond between two particular  atoms in a large number of compounds.  

 Factors that affect bond length

∙ Type of atoms

∙ Type of bond – single, double, or triple  

∙ Bond length of triple bonds < bond length of double bonds <  bond length of single bonds for a particular pair of atoms

∙ Bond length and bond energies – bond energy generally decreases as bond length increases, but it is not a smooth trend  



• Topic 1: VSEPR Theory

∙ Valence Shell Electron Repulsion Theory

 Theory states that bonding electrons and lone pair electrons repel each  other and this repulsion determines the molecular geometry  

∙ Definitions of important terms

 Electron groups – any lone pairs, single bonds, multiple bonds and single electrons on the central atom of the molecule; # e­ group = # LP + # BP

 Electron geometry – geometrical arrangement of the electron groups; determined by number of electron groups on central atoms; not affected by presence  of lone pairs

 Molecular geometry – geometrical arrangement of the atoms (actual shape of the  molecule); affected by TYPE of electron group – lone pair vs bonding pair  Ideal bond angle – the angle of each basic shape without lone pairs ∙ Five basic shapes (electron geometry)

 Linear Geometry – 2 electron groups

 Trigonal planar symmetry – 3 electron groups

 Tetrahedral geometry – 4 electron groups

 Trigonal bipyramidal geometry – 5 electron groups

 Octahedral – 6 electron groups

∙ Effect on lone pairs on the central atom

 Lone pairs (LP) take up more room than bonding pair (BP) electrons   LP, LP repulsion > LP, BP repulsion > BP, BP repulsion

 Repulsion between two lone pairs is highest

 Repulsion between two bonding pairs is lowest

 Because of the additional repulsion from the lone pair, molecules with lone pairs on the central atoms will have bond angles that deviate from the  ideal bond angle

• Topic 2: Molecular Geometry

∙ Molecular shapes  

 A = central atom; X = peripheral atom bonded to A, E = lone pair  electrons on A

 AX3E  

 three bonding pairs; one lone pair

 Tetrahedral electron geometry (b/c of 4 electron groups)

 LP, BP repulsion pushes atoms in the molecules close; therefore, the  actual angle is less than the predicted 109.5° for tetrahedral  

 Actual angle = 107°

 Molecular geometry = trigonal planar  

 Ex: NH3


 AX2E2

 two bonding pairs; two lone pairs

 Tetrahedral electron geometry (b/c of 4 electron groups)

 Two LP, BP repulsion pushes atom even closer than in trigonal  pyramidal molecule, which has one LP, BP repulsion

 Therefore, actual angle = 104.5°

 Molecular geometry = bent

 Ex: H2O

 AX4E

 four bonding pairs; one lone pair

 Trigonal bipyramidal electron geometry (b/c of 5 electron groups)  Two LP, BP repulsions exist 90° apart  

 Molecular geometry = seesaw

 Ex: SF4

 AX3E2

 Three bonding pairs; two lone pairs

 Trigonal bipyramidal electron geometry (b/c of 5 electron groups)  LP’s always go on equatorial positions b/c this position minimizes  repulsion

 Molecular geometry = t-shaped  

 AX2E3

 Two bonding pairs, three lone pairs

 Trigonal bipyramidal electron geometry (b/c of 5 electron groups)  LP’s always go on equatorial positions b/c this position minimizes  repulsion

 Actual angle = 180°

 Molecular geometry = linear

 AX5E

 Five bonding pairs, one lone pair

 Octahedral electron geometry (b/c of 6 electron groups)

 LP position does not matter because octahedral geometry is  symmetrical

 Molecular geometry = square pyramidal  

 AX4E2

 Four bonding pairs, two lone pairs

 Octahedral electron geometry (b/c of 6 electron groups)

 LP’s are as far away from each as can be (opposite positons) 42

 Molecular geometry = square planar

• Topic 3: Molecular Geometry and Polarity

∙ Nonpolar molecules with dipoles

 If molecular geometry causes the dipole movements to cancel each  other, the molecule will be nonpolar

 Ex: CO2

∙ How to determine if a molecule is polar

 Draw a Lewis structure and determine molecular geometry

 Determine if the molecule contains polar bonds

 Determine if the polar bonds add together to form a net dipole moment  Linear

∙ If a diatomic molecule has a dipole moment, the molecule will be  polar

 Bent

∙ Dipole moments do no cancel because net dipole exists even when  the two dipoles are pointing in opposite directions

 Trigonal planar

∙ If all three dipole moments point out, the dipoles will cancel and the  molecule will be nonpolar

∙ If there are less than three dipole moments, the molecule will be  polar

 Tetrahedral  

∙ If all four dipole moments point out, the dipoles will cancel and the  molecule will be nonpolar

∙ If there are less than four dipole moments, the molecule will be polar  Properties affected by polarity of molecules

∙ Polar molecules and non-polar molecules do not mix  

∙ Like dissolves like

∙ Dipole-dipole forces

 Intermolecular forces (between adjacent molecules)

∙ Occur between polar molecule because opposite partial charges  attract

∙ Dipole-dipole force is very weak

• Topic 4: Hybridization

∙ Relation to Valence Bond Theory

 Explains bonding using orbitals

 Atomic orbitals overlap to form a bond

 The greater the overlap between two orbitals, the stronger the bond  and the lower the energy of the molecule

 Hybrid orbitals create greater overlap with orbitals of other atoms  because electron probability density is more concentrated in a single  directional lobe


 Therefore, hybrid orbitals minimize the energy of the molecule by  maximizing the orbital overlap in a bond

∙ Types of bonds  

 Sigma – direct overlap between orbitals

 Pi – results from side ways overlap between orbitals

 Single bond = one sigma bond

 Double bond = one sigma bond, one pi bond

 Exists in sp2 bonding

 Triple bond = one sigma bond, two pi bonds

 Exists in sp bonding

∙ Hybridization

 Definition – a mathematical procedure that allows us to combine the  standard atomic orbitals to form new atomic orbitals that form HYBRID  ORBITALS that correspond more closely to the actual distribution of  electrons in chemically bonded atoms

 Purpose – to recognize that the orbitals in a molecule are not necessarily the  same as the orbitals in an atom

 Difference between standard orbitals and hybrid orbitals  

 hybrid orbitals are localized over individual atoms, just like in standard  orbitals.  

 Hybrid orbitals have different shapes and energies than standard  orbitals  

 General rules

o # of standard atomic orbital added together = number of hybrid  orbitals formed

o particular combinations of standard atomic orbitals added together  determine the shapes and energies of the hybrid orbitals formed o type of hybridization that occurs is the one that yields the lowest  overall energy for the molecules

o carbon tends to form 4 bonds in compounds, and therefore always  hybridizes

 sp3 hybridization

o tetrahedral geometry  

o “sp3” indicates that the hybrid orbitals are a mix of one s orbital and three p  orbitals

o all four of the hybrid orbitals have the same energy  

o each sp3 orbital contains one electron  

o ex: NH3 -> N is sp3 hybridized and there are 4 sp3 orbitals around N  sp2 hybridization

o trigonal planar geometry

o “sp2” indicates that the hybrid orbitals are a mix of one s orbital and two p  orbitals

o one unhybridized p orbital exists

o the unhybridized p orbital forms a double bond through sideways  overlap  

o ex: H2C


 sp hybridization  

o linear geometry

o hybrid orbitals are a mix of two s orbitals and two p orbitals

o two unhybridized p orbitals

o the two unhybridized p orbitals form two sideways overlap, creating a triple bond

o ex: N

 sp3d hybridization

o trigonal bipyramidal geometry  

o ex: PF5

 sp3d2 hybridization

o octahedral geometry  

o ex: SF6


• Topic 1: Oxidation Numbers

∙ For elements, oxidation number is 0

∙ For ions, the oxidation number = its charge

∙ For oxygen, the oxidation number is usually -2

∙ For hydrogen, the oxidation number is usually +1

∙ The sum of all the oxidation numbers = overall charge of a molecule or ion

• Topic 2: Redox Reactions

∙ Oxidation-Reduction Reactions

∙ Involves a change in oxidation number

∙ Oxidation = when reactant increases oxidation state

∙ Reduction = when reactant decreases oxidation state

∙ Redox reaction occurs when one reactant is oxidized and the other is  reduced

• Topic 3: Effective Nuclear Charge (Zeff)

∙ The charge that an electron “feels” due to shielding of the nucleus by the  inner electrons

∙ Zeff = Z - # of inner core e- 

∙ but this formula is not accurate

∙ Down a period, Zeff decreases

∙ Across a period, Zeff increases


∙ Across a period, radii decreases because of increased Zeff

• Topic 4: Melting Points

∙ While a compound melts, the temperature remains constant  ∙ When melting ice, break dipole-dipole forces between molecules ∙ The bond strength of an ionic compound is greater than the bond strength  of covalent compounds

∙ Therefore, MP of ionic compounds is much higher than MP of covalent  compounds

∙ As lattice enthalpy increases, MP increases

• Topic 5: Strong Acids

∙ Dissociate 100% in solution

∙ The strong acids are hydrochloric acid (HCl), hydrobromic acid (HBr),  hydroiodic acid (HI), nitric acid (HNO3), sulfuric acid (H2SO4), perchloric  acid (HClO4)

∙ On the other hand, weak acids are mainly undissociated in solution

• Topic 6: Molarity of Solutions and Ions

∙ Molarity = moles solute/ volume of solution

∙ Molarity = concentration



For Chapter 1, KNOW:

3 states of matter and how to represent them using  diagrams

Homogeneous vs heterogeneous mixtures

Crystalline vs amorphous solid

Physical and chemical properties

Basic SI units  

SI prefixes from tera through pico

How to do calculations using dimensional analysis/ unit  conversions involving mass, volume and density

Rules for significant figures

How to write and compute numbers in scientific notation

For Chapter 2, KNOW:

Laws of mass conservation

Law of definite proportions

Law of multiple proportions

How to do calculations with the above three laws Dalton’s Atomic Theory

Thomson’s, Millikan’s, Rutherford’s experiments Symbols for protons, neutrons, electrons  

Relative charges of protons, neutrons, electrons Atomic number vs mass number  

How to distinguish one isotope from the other  Cation vs anion

Mendeleev’s periodic law and periodic table

A group vs a period  

Location of metals, metalloids and nonmetals in the periodic  table

How to predict charges of ions

How to do calculation involving atomic masses and percent  abundance  

The concept of the mole and Avogadro’s number (NA) How to convert between moles and number of atoms How to convert between mass to moles using molar mass How to convert between mass and number of atoms

For Chapter 3, KNOW:

Difference between ionic bonding and covalent bonding How to predict if a compound is ionic or covalent

The differences between chemical formula, molecular,  empirical formula and structural formula

How to visualize molecules using structures, ball-and-stick  models and space-filling models

Atomic elements vs molecular elements  

How to go between chemical formula and chemical name  

How to calculate formula mass given chemical formula and  atomic masses

Mass percent composition and how it is calculated

How to determine the molecular formula and the empirical  formula from experimental data

How to balance chemical equation  

How to determine empirical formula of organic compounds  that are combusted when given the masses of CO2 and H2O Formulas of common anions and cations

Formulas of common polyatomic ions


For Chapter 4, KNOW:

How to go between moles or grams of a product and moles  or grams of a reactant when given a balanced chemical  equation

How to determine which is the limiting reactant and which is  the reactant in excess when given mole ratio of reactants in a  balanced chemical equation

How to calculate the moles or mass of one of the products  from the mass of each of the reactants

The limiting reactant is the one that yields the fewer moles  or lower mass of the product

The difference between theoretical yield of a product and the actual yield  

How to calculate the percentage yield

That if percent yield is greater than 100%, the product has  been contaminated or is wet  

Solute, solvent, solution

Dilute and concentrated solution

How to go between molarity and mass of solute and volume  of solution

How to use the dilution equation M1V1 = M2V2

How to calculate the molarity and volume of another  reactant given the molarity and volume of one reactant  

How to use the equation (M1V1)/n1 = (M2V2)/n2

Definition of precipitate  

Definition of strong, weak and non-electrolyte and how to  identify them

Symbol (aq)

How to write balanced molecular equation for a precipitation  reaction

How to predict a precipitate


How to write a full ionic equation

How to write a net ionic equation

How to identify spectator ions

Definition of acids and bases

Definition of neutralization

Definition of ionization  

Definition of polyprotic (acid)

How to write equation to show acids and bases dissolved in  water

The names and formulas of 6 strong acids and strong bases Difference between a strong acid and a weak acid Difference between a strong base and a weak base The double arrow ⇌ and what it means

Strong acids and bases are 100% ionized and are strong  electrolytes

Weak acid and bases are only partially ionized and are weak  electrolytes

The formula of a neutralization reaction is Acid + Base → Salt + Water

How to write a balanced molecular equation, a full ionic  equation and a net ionic equation for neutralization reactions

Use the equation (MaVa)/na = (MbVb)/nb for acid-base reactions in titrations

For Chapter 5, KNOW:

What pressure is  

What causes gas pressure

Various units of gas pressures

How to convert between various units of gas pressures 4

The equations of the Gas Laws: Boyle, Charles, Avogadro,  Amonton

How to do problems using Boyle’s, Charles’s Avogadro’s and  Amonton’s Laws

The equation of the Ideal Gas Law  

How to do problems using the Ideal Gas Law

How to use the equation (P1V1)/T1 = (P2V2)/T2 when n is  constant  

1 mole of an ideal gas occupies 22.4 L at STP

How to get density from the Ideal Gas Law through the  equation PM=dRT

Dalton’s Law of Partial Pressures

How to get mole fraction (X)and how to use X to get partial  pressures  

How to do calculations with Law of Partial Pressures

How to do stoichiometry problems and using moles to get  pressure and volume after getting the number of moles

The Kinetic Molecular Theory and its 3 postulates

How to use the Kinetic Molecular Theory to explain the  properties and behavior of gases

How to use the Kinetic Molecular Theory to explain the gas  laws  

The root mean square velocity, urms  

R in Urms equation is 8.314 J/(mol K) and MM in Urms is in  kg/mol

How velocity distributions vary depending on molar mass  and temperature

Difference between diffusion and effusion  

Graham’s Law of Effusion

Know that the Ideal Gas Law only holds true at moderate  pressures and temperatures  

For real gases, at high pressures, the size of the gas particles has to be taken into account


For real gases, at low temperatures, collisions are inelastic  and attractive forces have to be taken into account

What the van der Waal’s equation looks like and which  variable accounts for low temperature and which accounts for  high pressure

For Chapter 6, KNOW:

What is energy

How energy is related to work

The different types of energy  

What the internal energy of a system is

Different units of energy – J, kJ, cal, Cal

How to convert between units of energy

System and surroundings

Law of Conservation of Energy

First Law of Thermodynamics

How changes in internal energy relate to system and  surroundings

How to explain changes in internal energy with reference to  the system and its surroundings and the universe

Definition of a state function

Changes in internal energy is related to work and heat  The equation ΔE = q + w

q (heat) is negative when system releases energy and  positive when system absorbs energy

w (work) is negative when system does work on the  surroundings and positive when the surroundings do work on  the system

The equation q = mC ΔT  

The equation w = -P ΔV  

Expansion work is done by a system when it expands against an external pressure


How a bomb calorimeter works  

How to calculate qrxn and ΔErxn for a reaction done in a bomb  calorimeter  

Qcal = Ccal ΔT  

Heat gained by the calorimeter is equal to the heat lost by  the combustion reaction

+qcal = -qrxn

Enthalpy is related to internal energy and heat  

At constant pressure, ΔH = qp

At constant volume, ΔH = qv

How to calculate ΔHrxn from qrxn for reactions done in solution  in a coffee cup calorimeter using q = mCΔT

Heat gained by the solution equals the heat lost by the  reaction  

+qsoln = - qrxn

For Chapter 7, KNOW:

the differences types of EM radiation and what they are

Memorize the order in energy, wavelength and frequency of  the different types of EM radiation

Be able to do calculations using c = νλ

Planck’s equation E=hν = (hc)/λ

Electrons can jump from one level to another

Emission spectrum  

Absorption spectrum  

When an electron in the H atom jumps to level n=1, UV light  is emitted

When an electron in the H atom jumps to level n=2, visible  light is emitted

When an electron in the H atom jumps to level n=3, IR is  emitted


Rydberg equation 1/λ = R [1/m2 – 1/n2]

∆Eatom = -Ephoton

En = -2.18 x 10-18 J (1/n2)  

DeBroglie equation λ = h/(mν)

Heisenberg’s Uncertainty Principle

Schrödinger’s wave function, the probability of finding an  electron in an atom which describes the shape of the orbital

Quantum numbers n, l, ml , ms

Know the rules for possible combinations of quantum  numbers

Wave functions and probability give rise to orbitals

Orbitals are a mathematical representation of the probability  of finding an electron in that space

Know the shape of the orbitals s, px, py, pz and d

For Chapter 8, KNOW:

How to apply the aufbau principle

How to apply Pauli’s exclusion principle

How to apply Hund’s rule

How to write the expected electron configurations for  elements up to Radon

How to use the noble gas abbreviation  

The exceptional electron configurations for the transition  metals Cr and Cu

How to write the expected electron configuration of cations  and anions  

How to predicts charges for elements in groups 1, 2, 13-17  

Electrons are lost from the ns shell before the (n-1)d shell for  transition metals

Diamagnetic vs paramagnetic  

Trends for atomic radii across a period and down a group  8

Effective nuclear charge



Atomic radii vs ionic radii  

What is meant by isoelectronic

How to rank radii of isoelectronic ions using nuclear charges Ionization energy  

Trends in first ionization energy across the period and down  the group

Relative values of successive IEs depend on number of  valence electrons

How to identify an element given its successive IEs Electron affinity  

Trends in electron affinity across a period  

Noble gases (gp 18) have positive electron affinities  Metallic character trends across a period and down a group

For Chapter 9, KNOW:

Why chemical bonds form

how to draw Lewis dot structures of elements and ions in  group 1, 2, 13-18

Number of dots in a Lewis dot structure = number of valence electrons

The different types of chemical bonds – ionic, covalent,  metallic

How to draw simple Lewis dot structures of atoms How to represent ionic bonds using Lewis dot structures Definition of lattice energy

Relationship of lattice energy to strength of the ionic bond Trends in lattice energy

Ion size


Ion charge

Effect of lattice energy on other properties such as melting  point, boiling point and solubility

Octet rule

Duet rule

Bonding pair

Lone pair

Double and triple bonds

Electronegativity definition

Trend across a group

Trend down a group

How to predict the polarity of a bond using differences in  electronegativity  


Memorize the difference range that make a bond  nonpolar, polar and ionic

How to draw a Lewis dot structure for polyatomic ions  Definition of resonance

How to draw a resonant structure

Definition for formula charge

Formula for formal charge

What formal charges are most favorable  

Exceptions to the octet rule


Incomplete octets

Expanded octets

For Chapter 10, KNOW:

Electron geometry vs molecular geometry  


How to predict the electron geometry and molecular  geometry of a molecule given a Lewis dot structure that  shows all lone pair electrons around the atoms

How to recognize and draw the various molecular geometries on paper

Molecule’s shape = its molecular geometry

How to determine if a molecule has a net dipole moment  

A molecule is nonpolar if it has dipole moment that cancel  out or if there are no dipole moment

A molecule is polar if it has a net dipole moment

How to predict the hybridization of the atom depending of  the number and type of bonds around an atom  

The difference between sigma and pi bonds  

How to count the total number of each type of bonds in a  molecule


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