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CHEM 1030 FINAL EXAM Study Guide

by: Alyssa Anderson

CHEM 1030 FINAL EXAM Study Guide CHEM 1030

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Alyssa Anderson

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This is the final draft of the study guide for the chemistry final Wednesday, May 4, 2016. It includes all the information that will be on the final exam, starting at Week 1 and ending on Week 15 (...
Fundamentals Chemistry I
Dr. Streit
Study Guide
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Date Created: 04/25/16
1 CHEM 1030 FINAL EXAM STUDY GUIDE *Chemistry Exam 1* Chemistry- the study of matter and changes that matter undergoes Matter- anything that has mass and occupies space Scientific Method- a procedure/set of guidelines to organize and publish efficiently 1. Gather data through observations and experiments 2. Identify patterns and trends in collected data and note any initial thoughts 3. Summarize findings with a law- a concise statement that makes a relation between phenomena 4. Formulate a hypothesis by observing the cause and effect relationship 5. With time, the hypothesis may evolve into a theory which can predict future occurrences Classification of Matter 1. Solid- particles are held close together in an ordered position and DO NOT conform to the container it is placed in 2. Liquid- particles are held relatively close together but do not have an organized pattern and DO conform to the the container it is placed in 3. Gas- particles are far apart and have no set pattern but DO conform to the container it is placed in 4. In principle, all substances can exist in the solid, liquid, or gaseous stage 5. We can convert a substance by changing its identity 6. Mixtures can be separated by physical means into its component without changing the identities of the components - example: magnet to separate sand and iron (iron is magnetic) - example: boil water to separate salt and water (water has a much lower boiling point) - example: boil water to separate water and alcohol (different boiling points) 2 Chemists classify matter as either a substance or a mixture of substances 1. Substance- a form of matter that has a definite composition and distinct properties - example: salt (NaCl), iron, water (H20), mercury, carbon dioxide (CO2) - substances differ from each other in composition and may be identified by appearance, taste, smell, etc. 2. Mixture- physical combination of 2 or more substances a. homogenous- uniform throughout solution - example: seawater, apple juice, cake b. heterogenous- not uniform throughout solution - example: trail mix, chicken noodle soup, shells in sand Properties of Matter 1. Quantitative- properties measured/expressed with a number/unit (QUANTITY) 2. Qualitative- properties measured without measurements but rather are based on observations using the senses (taste, color, smell, etc.) (QUALITY) 3. Physical Property- one that can be observed or measured without changing the identity of the substance - example: color, melting point, boiling point 4. Chemical Property- one that a substance exhibits as it interact with another substance - example: flammability, corrosiveness, rust 5. Physical Change- change where the state of matter changes but the identity of the matter does not change - example: changes of state (melting, boiling, freezing, condensing) 6. Chemical Change- change in the composition so that the original composition no longer exists - example: digestion, combustion, oxidation 7. Extensive Property- depends on the amount of matter present - example: mass, volume, aka additive properties 8. Intensive Property- does NOT depend on the amount present - example: temperature, density 9. Physical Process- mixtures are separated but the identities do not change 10. Chemical Process- a process of changing mixtures/chemicals 3 Scientific Measurement 1. Properties that can be measured are called quantitative 2. A measured quantity must always include a unit 3. Systems A. English- foot, gallon, pound, Fahrenheit B. Metric- meter, liter, kilogram C. International System of Units (SI units)- universally used by scientists 1. Meter 2. Kilogram 3. Kelvin 4. Second 5. Ampere 6. Mole 7. Candela 4. Mass (g or kg or amu) A. a CONSTANT measure of amount of matter in an object/sample B. Gravity varies from location to location constant so weight = mass x gravity C. the mass of an atom is 1 amu= 1.6605378 x 10^-24 g 5. Temperature (Celsius or Kelvin or Fahrenheit) A. Celcius- for water, freezing point is O*C, boiling point 100*C B. Kelvin (*SI UNIT*)- “absolute” scale because 0 K is the absolute lowest C. K = *C + 273.15 OR C* = K - 273.15 D. Fahrenheit- for water, freezing point is 32*F and boiling point is 212*F E. *F = (9/5)(*C) + 32*F OR *C = (5/9)(*F - 32) 6. Volume (meter^3 or Liter) A. V = (length)^3 B. 1 dm^3 = 1 L C. 1 cm^3 = 1 mL 4 7. Density (kg/m^3) A. d = mass/volume so d = mass/length^3 B. solid = g/cm^3 C. liquid = g/mL D. gas = g/L E. example: if d1>d2 then m1<m2 OR v1>v2 Uncertainty in Measurements 1. Exact Numbers- defined values or counted numbers A. example: 1 kg = 1000 g B. example: 1 dozen = 12 items C example: 28 students in a class 2. Inexact Numbers- measured by anything but counting such as length, volume, mass A. It must be reported to indicate uncertainty by using significant digits B. The last digit reported is called the uncertain digit C. example: if we have an item against a ruler and we think it’s about 2.5 inches long, we know it’s for sure 2 inches but not sure about the .5, so we put 2.5 +/- 0.1 inch, and with a more accurate ruler we could put 2.45 inches +/- 0.01 inch D. Guidelines of Significant Figures 1. Any nonzero numbers ARE significant 2. Zeros between nonzero numbers ARE significant 3. Zeros to the LEFT of the first nonzero digit are NOT significant 4. Zeros to the RIGHT of the nonzero digits in decimals ARE significant 5. Zeros to the RIGHT of the last nonzero digit in a number without a decimal MAY OR MAY NOT be significant - example: 100 could have 1 2 or 3 significant figures 5 Calculations with Measured Numbers 1. Addition/Subtraction- line up the decimals and take the answer with the smaller amount of digits (rounding may be necessary) 2. Multiplication/Division- preform the action then take the fewer amount of digits from the original numbers given (rounding may be necessary) 3. NOTE: Be sure not to include exact numbers, such as the counted number - example: when finding the mass of each of 2 pennies, knowing together they equal 15 grams, 2 is not included in the measurement of significant figures. Therefor, since together they had 15 grams and that is 2 significant figures, your answer will have 2 significant figures 4. Rounding A. Leave rounding for the LAST step. DO NOT ROUND AFTER EACH STEP B. If the last digit is greater than 5, round UP (ex: 318.175 = 318.18) C. If the last digit is less than 5, round DOWN (ex: 318.174 = 318.17) 5. NOTE: Be aware of powers of 10. Make sure that you are calculating variables with the same power, then proceed 6. NOTE: Significant figures matter even when scientific notation changes 7. NOTE: Be sure to calculate the correct mass or volume before proceeding to find density or weight 8. Accuracy- how close the measurement is to the TRUE value 9. Precision- how close a series of measurements are to one another Using Units and Solving Problems 1. Conversion Factor- fraction in which same quantity is expressed one way in the numerator and another in the denominator - example: 1 inch = 2.54 cm aka 1 in/2.54 cm OR 2.54 cm/1 inch 2. Dimensional Analysis- use of conversion factors in problem solving A. Also known as the factor-label method B. example: convert 12 inches to meters (NOTE: only use significant figures of the thing you are converting (so 2 s.f. because 12 inches has 2 s.f.); 12 inches x 2.54 cm/1 inch x 1m/100 cm = 0.3042 = 30.30 m) 6 The Development of the Atom 1. An atom is the smallest quantity of matter that still retains the properties of matter 2. An element is a substance that cannot be broken down into two or more similar substances by any means (such as gold, oxygen, helium) 3. Atoms can also be divided smaller and smaller and eventually only a single atom remains. Dividing it further would make pieces that are no longer atoms. 4. Dalton- said atoms (of which matter consists of) are tiny, invisible particles. 5. Once a single atom has been obtained, dividing it smaller produces subatomic particles. 6. The nature, number, and arrangements of subatomic particles determine the properties of atoms 7. NOTE: LIKE charges repel each other, OPPOSITE charges attract 8. JJ Thompson (1856-1940)- noted easy were repelled by a plate with a negative charge and attracted to a plate bearing a positive charge. He concluded the rays were negatively charged. His contributions include: A. Proposed rays were actually a stream of negatively charged particles B. Negatively charged particles equaled electrons C. By varying the electric field and measuring the degree of deflection of cathode rays, Thompson determined the charge-mass ratio 9. R.A. Milikan (1868-1953)- determined the charge on an electron by examining the motion of tiny oil drops, which was found to be -1.6022 x 10^-19 C 10. The mass of an electron equals the charge divided by the charge multiplied by the mass which means (-1.6022 x 10^-19)/(-1.76 x 10^8 C x grams) which means it equals 9.10 x 10^-28 grams 11.Wilhelm Rotgen (1845-1923)- discovered x-rays which are not deflected by magnetic or electric fields so that they could not consist of charged particles 12. Antoine Becquerel (1852-1908)- discovered radioactivity 13. Alpha rays- consist of positively charged particles called alpha particles (α) 14. Beta rays- electrons that are deflected and made of beta particles (β) 7 15. Ernest Rutherford- used α particles to prove the structure of atoms A. The majority of particles penetrated the gold undeflected B. Sometimes, a gold particle would be deflected at a large angle or even backwards C. Through this, Rutherford concluded the nuclear model which states a positively charged center is concentrated in the middle of a cell at the nucleus and that the nucleus accounts for most of the cell’s mass and is extremely dense at the core within the atom 16. BE SURE TO LOOK AT TABLE 2.1/2.2 IN THE BOOK 17. All atoms can be identified by the number of protons/neutrons they have Characteristics of the Atom 1. Atomic Number- number of protons in the nucleus A. Since atoms must stay neutral, the number of protons equals electrons B. Protons determine the identity of the element 2. Mass Number- number of protons added to neutrons 3. Isotopes- Most atoms have at least two, which mean they have the same amount of protons and electrons, but different number of neutrons, which effects the mass. They usually exhibit the same chemical properties, such as some have the same type of compound with similar reactivities. On occasion, an isotope will be radioactive. 4. Nuclear Stability- can be related to density (note: the total volume is hardly accounted for by the nucleus but the mass is mainly the nucleus alone) A. The higher the density, the stronger the forces are in the atom. B. Stability = Coulomb repulsion - short range attraction C. example: the atomic number of Uranium equals 92 but has 143 neutrons D. Heavy atoms need much more neutrons to remain stable E. The principle factor for nuclear stability is proton to neutron ratio (n/p) F. There are more stable nuclei with 2, 8, 20, 50, 82, or 126 protons and neutrons G. There are more stable nuclei with even numbers H. All elements with atomic numbers greater than 83 are radioactive 6. Atomic Mass- mass of atom in amu (1 amu = half the mass of a carbon-12 atom) 7. Average atomic mass- on the periodic table, it represents the average mass of the naturally occurring mixture of isotopes 8 Elements of the Atom . 1 Protons- positively charged, in the nucleus 2. Neutrons- no charge, in nucleus, slightly larger than protons 3. Electrons- negatively charged particles that orbit around a nucleus The Periodic Table 1. A chart in which elements having chemical and physical properties are grouped together, separated by periods and groups. 2. Numbered by increasing atomic number (protons and electrons) because it regulates all properties of that element (fingerprint of element) 3. There are two important numbers- the average atomic mass and the atomic number 4. Periods- horizontal rows, in order of increasing atomic number A. Metals- good conductors of heat and electricity B. Nonmetals- poor conductors of heat and electricity C. Metalloids- intermediate properties 5. Groups- vertical columns A. Alkali Metals (1A)- Li, Na, K, Rb, Cs, Fr B. Alkaline Earth Metals (2A)- Be, Mg, Ca, Sr, Ba, Ra C. Chalcogens (6A)- O, S, Se, Te, Po D. Halogens (7A)- F, Cl, Br, I, At E. Noble Gases (8A)- He, Ne, Ar, Kr, Xe, Rn F.Transition Metals (1B and 3B-8B) Mole- the amount of a substance that contains as many elementary entities as there are atoms in exactly 12 grams of carbon-12 1. Experimentally determined number- Avagandro’s Number (N ) A 2. NAaka 1 mole = 6.0221415 x 10^23 (usually simplified to 6.022 x 10^23) 3. example: The human body has a total of 30 moles of calcium. Determine the number of atoms of calcium and the number of moles of Calcium in a sample containing 1.00 x 10^26 Ca atoms. work: (30 moles Ca) x (6.022 x10^23/ 1 mole Ca) = 1.807 x 10^25 atoms of Ca 9 Molar Mass- mass in grams of 1 mol of substance 1. By definition, the mass of one mole of carbon-12 is exactly 12 grams 2. The mass of 1 carbon-12 atom equals exactly 12 amu 3. The mass of an atom equals the mass of the mole (just in different units) 4. example: determine the number of moles of carbon in 25 grams of carbon work: (25 grams of C) x (1 mole of C/ 12.01 grams of C) = 2.082 moles of C (10.50 grams He) x (1 mole He/ 4.003 grams C) = 2.633 mol He 5. example: determine the number of moles of He in 10.50 grams of helium work: (0.515 g C) x (1 mol C/ 12.01 g C) x (6.022 x 10^23/ 1 mole C) = 2.58 x 10^22 C atoms Unit of Energy 1. Measured by the Joule (J) 2. Created by English physicist James Joule 3. The amount of energy possessed by a 2 kg mass moving at speed of 1 m/s 4. E K 1/2 x m x u^2 = 1/2 x 2 kg x (1 m/s)^2 = 1 kg x m^2/s^2 5. Joules can also be denied as the amount of energy exerted when a force of 1 Newton is applied over 1 meter. 1 J = 1 N x m The nature of light- visible light is only a small component of the continuum of radiant energy known as the electromagnetic spectrum Properties of waves (all forms of electromagnetic radiation travel in waves) 1. Waves are characterized by wavelength (λ) which is the distance between identical points on successive waves (inversely related to v) 2. frequency (v, nu) is the number of waves that pass through a particular point in 1 second 3. NOTE: short wavelength = high frequency; long wavelength = low frequency 4. Amplitude is the vertical distance of the middle of the wave 5. The speed of light (c) through a vacuum is constant. c = 2.99792458 x 10^8 m/s 6. The speed of light, frequency, and wavelength are all related by the equation c = λ x v (λ is expressed in meters, and v is expressed in s^-1) 10 Quantum Theory 1. Atoms and photons at the microscopic level do not measure equally to the microscopic level 2. The laws for macroscopic are not applicable to the microscopic level 3. All energy is transferred through the measure of waves (E = h x v) 4. E = h x v means energy is calculated by multiplying Planks Constant (6.63 x 10^-34 J x s) by the frequency The Schrodinger Equation 1. Erwin Schrodinger realized the wave and particle characteristics were different in electrons 2. Particle behavior is determined by mass (m) while wave behavior is determined by the wave function (Ψ) in the equation H (m) x Ψ = E x Ψ 3. Quantum Mechanics- defines the region where the electron is most likely to be at a given time 4. The probability of finding an electron in a certain area of space is proportional to Ψ^2 and is called electron density 5. Energy states and wave functions are characterized by a set of quantum numbers 6. Quantum numbers and wave functions describe atomic orbitals Atomic Orbitals 1. All s orbitals are spherical in shape but alter in size ( 1s < 2s < 3s) 2. All p orbitals are dumbbell shaped and have 3 orientations 3. D orbitals vary and have 5 orientations 5. F orbitals vary and have 7 orientations 6. Energy of orbitals- in a hydrogen atom, depends only on n SUMMARY 1. Principle (n)- SIZE 2. Angular (l )- SLOPE/SHAPE 3. Magnetic (ml )- ORIENTATION 11 Quantum Numbers 1. They are required to describe the distribution of electron density in an atom 2. In order to describe an atomic orbital, you must know the three quantum numbers 3. Principal Quantum Number (n) A. Designates SIZE of the orbital B. The larger the value of n the larger the orbitals C. The allowed values of n are integral numbers (1, 2, 3, etc.) D. The collection of orbitals with the sam value of n are frequently called shells 4. Angular Momentum Quantum Numbers (l ) A. Describes the SHAPE of the orbital B. Values of l are integers that depend on the value of n C. Allowed values of l range from 0 to n-1 D. The collection is called a subshell 5. Magnetic Quantum Number (ml ) A. Determines the ORIENTATION of orbitals in space B. Values of ml are integers that depend on the value of l C. - l , 0, + l 6. Quantum Numbers designate shells, subshells, and orbitals- REFER TO TABLE 3.2 Speed (ms) 1. It is not derived by the equation 2. Found through experiments that included a beam of atoms that were split by a magnetic field 3. It was concluded that electrons behave like tiny magnets 4. Specifies the electrons spin 5. ms = +/- 1/2 Example: 2p^2 1. n = 2 2. l = p = 1 3. m l = +1, 0, -1 4. ms = +1/2, -1/2 12 Aufbau Principle 1. States that electrons are added to the lowest energy orbitals first before moving to higher energy orbitals 2. Example: Li has 3 electrons so its configuration is 1s^2 / 2s^1 3. Example: Be has 4 electrons so its configuration is 1s^2 / 2s^2 4. Example: B has 5 electrons so its configuration is 1s^2 / 2s^2 / 2p^1 5. Example: C has 6 electrons so its configuration is 1s^2 / 2s^2 / 2p^2 6. Example: F has 9 electrons so its configuration is 1s^2 / 2s^2 / 2p^5 7. NOTE: 2p orbitals are degenerate Pauli Exclusion Principle 1. No two electrons in an atom can have the same four quantum numbers 2. The principle number, angular momentum number, magnetic number, and speed cannot ALL be the same 3. Only 2 electrons can occupy an atomic orbital Hund’s Rule 1. The most stable arrangement of electrons is the one in which the number of electrons with the same spin is maximized 2. In other words, put 1 electron in each box before pairing NOTE: All the chemical and physical properties of matter are given by how the electrons are arranged in each orbital. Paramagnetism is when there are one or more unpaired electrons in an atom (such as the case of O and F). Diamagnetism is when all the electrons in an atom are paired, such as Neon. 13 Electron Configuration 1. Describes where the electrons are distributed in the various atomic orbitals 2. In the ground state of hydrogen, the electron is found in the 1s orbital (1s^1 means the principal number n = 1 and the angular momentum is s = 0). If hydrogens electrons were found in a higher energy we would say the atom is in an excited state (2s^1) 3. In multi-electron atoms, the orientations of the orbitals are SPLIT (i.e. goes from 3s to 3p then from 4s to 3d) Rules of Electron Configuration 1. Electrons will reside in the available orbitals of the lowest possible energy 2. Each orbital can accommodate a maximum of two electrons 3. Electrons will not pair in degenerate orbitals if an empty orbital is available 4. Orbitals will fill in the order indicated in the figure to the right. Worked example 3.10 Problem: What’s the electron configuration and orbital diagram of Ca (Z= 20)? Solution: 1s^2 / 2s^2 / 2p^6 / 3s^2 / 3p^6 / 4s^2 Noble Gas Core 1. The electron configurations of all elements except H and He can be represented by using a noble gas core 2. K (Z =19) has the configuration 1s^2 / 2s^2 / 2p^6 / 3s^2 / 3p^6 / 4s^1 but since argon (Ar) is 1s^2 / 2s^2 / 2p^6 / 3s^2 / 3p^6 you can adjust and only write [Ar] 4s^1 14 Electron Configuration and the Periodic Table 1. Valence electrons are the outer electron involved in chemical reactions and can be identified by the period number 2. 4f = the lanthanide (rare earth) series 3. 5f = the actinide series 4. Notable exceptions to electron filling in the transition metals: A. Chromium (Z = 24) is [Ar] s3^1 / 3d^5 B. Copper (z = 29) is [Ar] 4s^1 / 3d^10 C. The reason for these anomalies is the slightly greater stability of d subshells that are either half filled (d^5) or completely filled (d^10) Discoveries in the Periodic Table A. In 1864 John Newlands noted that when the elements were arranged in order of atomic number, every eighth element had similar properties. They could be grouped according to their properties and he called it the law of octaves. B. In 1869 Dmitri Mendeleev and Lothar Meyer independently proposed they idea of periodicity. 1. Mendeleev grouped the 66 known elements according to their properties and atomic mass 2. Mendeleev predicted properties for elements not yet discovered such as gallium (Ga) 3. However, Mendeleev could not explain inconsistencies such as argon coming before potassium in the periodic table despite having a higher atomic mass 4. In 1913 Henry Mosley discovered the correlation between the number of protons (atomic number) and frequency of x-rays generated. 5. By ordering the periodic table by atomic number instead of atomic mass, scientist were able to make sense of discrepancies. 6. Entries today include atomic number and symbol and are arranged according to electron configuration 15 Effective nuclear charge (Z subscript eff) is the actual magnitude of positive charge that is ”experienced” by an electron in the atom 1. In a multi electron atom, electrons are simultaneously attracted to the nucleus and repelled by one another (positive nucleus and negative electrons) 2. This results in shielding, where an electron is partially shielded from the positive charge of the nucleus by the other electrons 3. Although all electrons shield one another to some extent, the most effective are the core electrons 4. In general, the effective nuclear charge is given by Z , which is the number of protons in the nucleus 5. σ is the shielding constant Effective nuclear charge 16 Atomic radius is the distance between the nucleus of an atom and its valence shell 1. Atomic radius in metals aka metallic radius is half the distance between the nucleus of two adjacent identical metal atoms 2. Atomic radius in nonmetals aka covalent radius is half the stance between adjacent identical nuclei connected by a chemical bond 3. When we add another layer (shell) of electrons, we increase the radius. Therefore, the radius increases going down the periodic table. 4. When we go across a period, we are not changing the principal quantum number (n) but we are adding an angular number (l ). When l increases, the attraction between the effective nuclear charge and the charge on the valence shell becomes stronger so the electrons are pulled closer in (radius decreases). . 5 The atomic radius decreases left to right across a period due to the increased electrostatic attraction between the effective nuclear charge and the charge on the valence shell. 6. NOTE: The stronger the charge = the higher the attraction = the smaller the radius Worked example: Problem: referring only to a periodic table arrange the elements P, S, O in order of increasing atomic radius. Answer: P, S, O 17 For more help on the periodic table check out! 18 *Chemistry Exam 2* Ionization Energy- minimum energy required to remove an electron from an atom in a gas phase, resulting in an ion (chemical species with net charge POSITIVE charge means it is a CATION A. Example: Na (g) —> Na (g) + e - B. Therefore, sodium has an ionization energy of 495.8 kJ/mol (first ionization energy of sodium). C. IE1(Na) which corresponds to the removal of the most loosely held electrons In general, as Z effincreases, ionization energy also increases Higher atomic radius = lower Z eff= lower ionization energy Lower energy = closer to the radius = more stable NOTE: removing a paired electron is easier because of the repulsion forces between 2 electrons NOTE: Removing electrons can lower energy IE 1alues for main group elements (kJ/mol) are to the right It is possible to remove additional electrons in subsequent ionizations, giving IE1, IE2,tc. It takes more energy to remove the 2nd/3rd/4th electrons because it’s harder to remove core electrons than valence electrons 19 IE1(Mg) > IE (N1) because Mg is to the right so if it has a greater Z effnd more difficult to remove (496 kJ/mol < 738 kJ/mol) IE2(Na) > IE (M2) because the second ionization of Mg removes a valence electron where the second ionization of Na removes a core electron Electron affinity (EA) is the energy released when an atom in the gas phase accepts an electron Cl (g) + e —> Cl (g) Like ionization energy, electron affinity increases from left to right across a period as Z effincreases NOTE: It’s easier to add an electron to an empty orbital than to add an electron into an full orbital More than 1 electron may be added to an atom. O (g) + e^- —> O^ - (g) (EA = -141 kJ/mol) 1 O^ - + e^ - —> O^ 2- (g) (EA = -721 kJ/mol) While many 1st electron affinities are positive, subsequent EA’s are ALWAYS NEGATIVE. Considerable forces must be used to overpower compulsion energy
 EA 1Si) > EA (A1) because Si is more to the right therefore it has a greater Z eff EA 1Si) > EA (P1 because even though P is more to the right it involves putting it in a 3p orbital. The energy it costs of PAIRING outweighs advantage of adding electrons to atoms with easier Z eff 20 Metallic Character A. Metals tend to be shiny, lustrous, malleable, ductile, and good conductors of heat and electricity B. They have low IE (because they have CATIONS and not ANIONS) C. Many of the periodic trends of elements can be explained using Coulombs Law which states that the force (F) between two charged objects (Q1 and Q2) is directly proportional to the product of the two charges and inversely proportional to the distance (d) between the objects squared Ions of main group element A. Species with identical electron configurations to the noble gas to the right are called isoelectronic B. Common monatomic ions are arranged by their positions in the periodic table C. Exceptions: Mercury is actually a polyatomic ion (Hg ^2+) Electron Configuration of Ions A. Write the original configuration B. Add/remove to make appropriate number of electrons C. Note: in CATIONS, electrons are removed from the occupied orbitals with the highest value of n, but in ANIONS, electrons are added to the empty/partially filled orbitals with the lowest value of n NOTE: A. Low electron affinity- hard to accept electrons B. High electron affinity- easy to accept electrons C. Low ionization energy- easy to lose electrons D. High ionization energy- hard to lose electrons E. Noble gases have lower ionization energies 21 Ions of d-block elements A. Ions of d block elements are formed by removing electrons first from the shell with the highest value of n. (When you're removing electrons from the final configuration, you’re making it more stable by lessening the repulsion forces.) B. Always go to fill s first, then p. C. Exception: when you're trying to fill the p blocks Worked example 4.8: write the electron configuration for the following ions of d block elements A. Zn^2+ —> [Ar] 3d^10 B. Mn ^2+ —> [Ar] 3d^5 C. Cr ^3+ —> [Ar] 3d^5 Ionic radius- the radius of a cation or an anion (radius- distance between valence shell and nucleus) A. When an atom loses an electron to become a CATION, its radius decreases due in part to a reduction on electron-electron repulsions in the valence shell B. A significant decreases in radius occurs when ALL of an atoms valence electrons are removed C. When an atom gains one or more electrons and becomes an ANION, its radius increases due to increased electron-electron repulsions D. When referring to radius, a cation < neutral atom < anion Isoelectronic Series A. An isoelectronic series is a series of two or more species that have identical electron configurations but different nuclear charges B. Ionic radius and attraction force are difference C. Example: O^2- and F^- have the same electron configuration but different attraction forces NOTE: Stronger attraction forces make ionic radius shrink Worked example 4.9- know how to rank them in order of increasing radius 22 A compound is a substance composed of two or more elements combined in a specific ratio and held together by chemical bonds, such as salt (NaCl) or water (H20) Lewis dot symbols A. When atoms form compounds, their valence electrons actually interact B. A LDS consists of the elements symbol with dots (do not pair until needed) C. Examples 1. Boron —> 1s2/2s2/2p1 = 3 valence electrons 2. Carbon—> 4 valence electrons 3. Nitrogen—> 5 valence electrons D. The atoms combine in order to achieve a more stable electron configuration E. Maximum stability results when a chemical specifies is isoelectronic with a noble gas For main group metals such as Na the number of dots is the number of electrons that are lost For nonmetals in the second period the number of unpaired dots is the number of bonds the atom can form Ions may also be represented by Lewis dot symbols (O^2-) which can be represented by writing the element with correct number of electrons then brackets with the charge on the outside Ionic Compound Bonding- refers to the electrostatic attraction that holds oppositely charged ions together in an ionic compound A. Metals- low ionization energies lose electrons very easily B. Nonmetals- high electron affinity accepts electrons very easily C. Ionic Compounds 1. Lewis Dot Symbol 2. Na (dot) —> Na + e - 23 Lattice Energy A. A 3-dimensional array of oppositely-charged ions is called a lattice B. Lattice energyis the amount of energy required to convert a mole of ionic solid to its constituent ions in the gas phase C. The magnitude of lattice energy is a measure of an ionic compounds stability D. Lattice energy depends on magnitudes of the charge and on the distance between them E. If they have the same charges, the only thing they have that differs one from another is the distance between them F. *Know how to arrange ionic compounds in order of increasing/decreasing lattice energy (opposite of atomic radius)* G. NOTE: big atomic radius means low lattice energy H. The formation of ionic bonds RELEASES a large amount of energy The resulting electrically neutral compound, sodium chloride, is represented with the chemical formula NaCl 24 Naming Ions and Ionic Compounds A. A monoatomic cation is named by adding the word ION to the name of the element B. A monoatomic anion is named by changing the ending of the elements name to IDE, such as oxide, carbide, sulfide C. Some metals can form cations of more than one possible charge (especially true for elements of D-block) 2+ 1. Fe : ferrous ion [Fe(II)] 3+ 2. Fe : ferric ion [Fe(III)] 3. Mn : manganese (II) ion 4. Mn : manganese (III) ion 5. Mn : manganese (IV) ion Formulas of Ionic Compounds A. Ionic compounds are electronegatively neutral 1. In order for ionic compounds to be electrically neutral, the sum of the charges of the cation and anion in each formula must be zero 2. Example: aluminum oxide. 2(+3) + 3(-2) = 0 B. To name ions and ionic compounds: 1. Name the cation by omitting the word ion and using a roman numeral if the cation can have more than one charge 2. Name the anion by adding the word IDE 3. Examples a. NaBr —> Sodium Bromide b. CaO —> Calcium Oxide c. Mg 3 —2 Manganese Nitride d. Fe 2 3> Iron(III) Sulfide C. NOTE: the subscript of the anion is the charge of the cation D. NOTE: the subscript of the cation is the charge of the anion 25 Covalent Bonding and Molecules A. When compounds form between elements with similar properties with similar properties, electrons are not transferred from one electron to another but instead are shared in order to give each atoms a noble gas configuration B. This approach is known as the Lewis Theory of Bonding, named for its proponent, Gilbert Lewis C. Lewis Theory depicts bond formation in H2 (right) D. A molecule may be an element or a compound E. Different samples of a given compound always contain the same ratio, known as the law of definite proportions F. Amolecule is a combination of at least two atoms in specific arrangement held together by chemical forces (chemical bonds) G. A molecule may be an element or a compound H. Different samples of a given compound always contain the same ratio. This is known as the law of definite proportions. I. If two elements can form two or more different compounds, the law of multiple proportions tells us that when the masses of two elements when combined with each other to form more than one compound are in a ratio of small whole numbers J. The mass ratio of oxygen to carbon dioxide is 2.66:1, and the ratio of oxygen to carbon in carbon monoxide is 1.33:1 K. The ratio of two such mass rations can be expressed as small whole numbers L. Diatomic molecules contain two atoms and may either be heteronuclear or homonuclear (a) would be homonuclear diatomic but it has more than two so it is not diatomic. Imagine the same colors but with only two atoms bonded. (b) is polyatomic (c) is heteronuclear diatomic 26 Formulas A. A chemical formula denotes the composition of the substance B. A molecular formula shows the exact number of atoms in each element in a molecule C. Some elements have two or more distant forms known as allotropes, such as oxygen (O2) and ozone (O3) are allotropes of oxygen D. A structural formula shows not only the elemental composition but also the general arrangements E. An empirical formula uses whole-number ratios of elements to get the smallest bit of information that we can get from observation. While the molecular formulas tell us the actual number of atoms (the true formula), the empirical formula gives the simplest formula. Sometimes the true formula IS the empirical formula, like with H2O. 1. Molecular formula: N2H2 2. Empirical formula: NH2 3. Look at the worked example 5.6 Naming Molecular Compounds A. Remember: binary molecular compounds are substances that consist of just two different elements B. Nomenclature 1. Name the first element that appears in the formula 2. Name the second element that appears in the formula, change ending to IDE C. Greek prefixes are used to denote the number of atoms of each element 1. mono- is usually omitted for the first element 2. For ease of pronunciation, we usually eliminate the last letter of a prefix that ends in “o” or “a” when naming an oxide 3. Worked example a. NF3 nitrogen trifluoride b. N2O4 dinitrogen tetroxide 27 D. Compounds containing hydrogen 1. The names of molecular compounds containing hydrogen do not usually conform to the systematic nomenclature guidelines 2. Many are called by the common, nonsystematic names or by names that do not indicate explicitly the number of H atoms present a. B2H6 Diborane b. SiH Silane 4 c. NH3 Ammonia d. PH3 Phosphine e. H2O water f. 2 S hydrogen sulfide E. One definition of an acid is a substance that produces hydrogen ions (H ) when dissolved in water (HCl is an example of a binary compound that is an acid when dissolved in water) 1. To name these acids: a. Remove the –gen ending from hydrogen b. Change the –ide ending to on the second element to –ic (Hydrogen chloride hydrochloric acid) 2. A compound must contain at least one ionizable hydrogen atom to be an acid upon dissolving Organic Compounds A. Our nomenclature discussion so far as focused on inorganic compounds, generally defined as those without carbon B. Organic compounds contain carbon and hydrogen, sometimes in combination with other atoms C. Hydrocarbons contain only carbon and hydrogen D. The simplest hydrocarbons are called alkanes E. Many organic compounds contain groups of atoms known as functional groups, which often determine a molecule’s reactivity 28 Covalent Bonding in Ionic Species A. Polyatomic ions consist of a combination of two or more atoms B. Formulas are determined following the same rule as for ionic compounds containing only monatomic ions: ions must combine in a ratio that gives a neutral formula overall C. Oxoanions are polyatomic anions that contain one or more oxygen atoms and one atom (the central atom) of another element D. Starting with the oxoanions that end in –ate, we can name these ions: 1. The ion with one MORE O atom atom than the –ate ion is called the per…ate ion. (ie ClO - 3 a chlorate ion, so CLO - = 4 perchlorate ion) 2. The ion with one LESS O atom than the -ate ion is called the –ite ion. (ie ClO2- is the chlorite ion) 3. The in with TWO FEWER O atom than the -ate ion is called the hypo…ite ion. (ie ClO- is the hypochlorite ion) E. Oxoanions 1. percholate ClO 4 2. chlorate ClO - 3 3. chlorite ClO - 2 4. hypochlorite ClO- 5. nitrate NO - 3 6. nitrite NO -2 7. phosphate PO 43- 8. phosphite PO 33- 9. sulfate SO 42 10. sulfite SO 32- F.Oxoacids, when dissolved in water, produce hydrogen ions and the corresponding oxoanions 1. An–ate ion is called ……. ic acid (HClO chloric acid) 3 2. An –ite ion is called …….. ous acid (HClO chlo2ous acid) 3. Prefixes in oxoanions names are retained in naming oxoacids 4. Oxoacids, can be monoprotic (one ionizable hydrogen) or polyprotic (more than one ionizable hydrogen) 29 Hydrates A. A compound that has a specific number of water molecules within its solid structure B. For example, in its normal state, copper(II) sulfate has five water molecules associated with it C. Copper(II) sulfate pentahydrate —> CuSO4 x 2H2O D. When the water molecules are driven off by heating, the resulting compound, Cu(SO)4 is sometimes called anhydrous copper(II) sulfate E. Anhydrous means the compound no longer has water molecules associated with it F. Cu(SO)4 is white but Cu(SO)4 x 5H2O is blue Molecular and Formula Mass A. The molecular mass is the mass in atomic mass units (amu) of an individual molecule B. To calculate the molecular mass, multiply the atomic mass for each element in a molecule by the number of atoms of that element and total the masses C. Molecular mass of H2O = 2(atomic mass of H) + atomic mass of 0 D. Although the ionic compound does not have a molecular mass, we can use its empirical formula to calculate its formula mass E. Because the atomic masses on the periodic table are average atomic masses, the result of such a determination is an average molecular mass, sometimes referred to as the molecular weight. Percent Composition of Compounds A. A list of the percent by mass of each element in a compound is known as the compounds percent composition by mass B. Percent mass of an element = (n x atomic mass of element) / (molecular or formula mass of compound) x 100% C. We could have also used the empirical formula of hydrogen peroxide (HO) for the calculation D. In this case, we could have used the empirical formula mass to find the mass in amu of one of the compounds E. Worked Example 5.13 30 Molar mass (M ) A. In a substance, the molar mass is the mass in grams of one mole of the substance B. The molar mass of an element is numerically equal to its atomic mass 1. 1 mol C = 12.01 g 2. 1 C atom = 12.01 amu C. The molar mass of a compound us the sum of the molar masses of the elements it contains: 1. 1 mol H2O = 2 x 1.008 g + 16g = 18.02 g 2. 1 mol NaCl = 22.99 g + 35.45g = 58.44g The Octet Rule A. According to the octet rule, atoms will lose, gain, or share electrons in order to achieve a noble gas electron configuration B. Only valence electrons contribute to bonding C. Only two valence electrons participate in the formation of the F2 bond Lewis Structure A. A lewis structure is a representation of covalent bonding B. Shared electron pairs are shown either as dashes or as pairs of dots C. Lone pairs are shown as pairs of dots on individual atoms D. In a single bond, atoms are held together by one electron pair E. In a double bond, atoms share two pairs of electrons F. In triple bond, atoms are held together by three electron pairs G. Bond length is defined as the distance between the nuclei of two covalently bonded atoms H. Multiple bonds are shorter than single bonds I. For a given pair of atoms, triple bonds are shorter than double bonds which are shorter than single bonds J. We quantify bond strength by measuring the quantity of energy required to break it 31 Electronegativity and Polarity A. There are two extremes in the spectrum of bonding: 1. Covalent bonds occur between atoms that SHARE electrons 2. Ionic bonds occur between a metal and nonmetal and involve ions B. Bonds that fall between these extremes are polar C. In polar covalent bonds, electrons are shared but not shared equally (the delta is used to denote partial charges on the atoms) E. Pure covalent bonds- neutral atoms held together by equally shared electrons G. Ionic bonds- oppositely charged ions held together by electrostatic attraction H. Electron density maps show the distributions of charge I. Typically, electrons spend a lot of time in red and very little time in blue Electronegativity A. Electronegativity is the ability of an atom in a compound to draw electrons to itself B. There is no sharp distinction between non-polar covalent and polar covalent or between polar covalent and ionic C. The following rules help distinguish among them: 1. A bond between atoms whose electronegativities differ by less than 0.5 is generally considered purely covalent or nonpolar 2. A bond between atoms who's electronegativities differ by the range of 0.5 to 2.0 is generally considered polar covalent 3. A bond between atoms whose electronegativities differ by 2.0 or more is generally considered ionic 32 Dipole Movement, Partial Charges, and Percent Ionic Character A. Regions where electrons spend little time is typically blue while regions where electrons spend a lot of time is tropically red B. An arrow is used to indicate the direction of electron shift C. A quantitative measure of the polarity of a bond is its dipole movement (mu) D. mu = Q x r (Q is the charge; r is the distance between the charges, and mu is always positive and expressed in debye units denoted by the letter D) E. 1 D = 3.336 × 10 -30coulomb meter. F. Although the designations “covalent,” “polar covalent,” and “ionic” can be useful, sometimes chemists wish to describe and compare chemical bonds with more precision. G. Comparing the calculated dipole moment with the measured values gives us a quantitative way to describe the nature of a bond using the term percent ionic character. Lewis Structures and Formal Charge A. Formal charge can be used to determine the most plausible Lewis Structure when more than one possibility exists for a compound. B. To determine associated electrons: 1. All the atom’s nonbonding electrons are associated with the atom 2. Half the atom’s bonding electrons are associated with the atom C. When there is more than one possible structure, the best arrangement is determined by the following guidelines: 1. A Lewis structure in which all formal charges are zero is preferred 2. Small formal charges are preferred to large formal charges 3. Formal charges are associated with electronegativities. Resonance A. A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure B. Resonance structures are a human invention C. Resonance structures differ only in the positions of their electrons 33 Exceptions to the Octet Rule: Incomplete Octets A. The central atom has fewer than eight electrons due to a shortage of electrons 1. Elements in group 3A also tend to form compounds surrounded by fewer than eight electrons 2. Boron, for example, reacts with halogens to form compounds of the general formula BX3 having six electrons around the boron atom B. The central atom has fewer than eight electrons due to an odd number of electrons. (Molecules with an odd number of electrons are sometimes referred to as free radicals) C. The central atom has more than eight electrons 1. Atoms in and beyond the third period can have more than 8 valence e’s 2. In addition to the 3s and 3p orbitals, elements in the third period also have 3d orbitals than can be used in bonding D. The bond between B-N has both the electrons contributed by the N atom 1. This type of bond is a coordinate covalent or dative bond 2. This type of bond formation is an example of a Lewis acid-base process BF 3s a Lewis acid: it can accept a pair of electrons NH 3s a Lewis Base: it donates a pair of electrons 34 WORKED EXAMPLES FOR EXAM 2 Worked Example 5.2 Arrange MgO, CaO, and SrO in order of increasing lattice energy. Consider the charges on the ions and the distances between them. Apply Coulomb’s law to determine the relative lattice energies. All three compounds contain and all three cations are +2. Recalling that lattice energy increases as the distance between ions decreases, we need only consider the radii of the cations as all three contain the same anion. From Figure 4.13, the ionic radii are 0.72 Å (M), 1.00 Å (Ca2+), and 1.18 Å (Sr+). Worked example 5.3 Name the following ionic compounds: (a) CaO, (b) Mg3 2, and (c) Fe2 3. Begin by identifying the cation and anion in each compound, and then combine the names for each, eliminating the word ion. Solution (a) CaO is (b) Mg3 2is (c) Fe2 3is Think About It Be careful not to confuse the subscript in the formula with the charge in the metal ion. In part (c), for example, the subscript on Fe is 2, but this is an iron(III) compound. 35 Worked Example 5.6 Write the empirical formulas for the following molecules: (a) glu6 12 6 ), a substance known as blood sugar; (b) adenine (5 5 5 ), also known as vitamin4B ; and (c) nitrous oxide (N2O), a gas that is used as an anesthetic (“laughing gas”) and as an aerosol propellant for whipped cream. Worked Example 5.7 Name the following binary molecular compounds: (a) N3 and (b) 2 4 . Strategy Each compound will be named using the systematic nomenclature including, where necessary, appropriate Greek prefixes. Think About It Make sure that the prefixes match the subscripts in the molecular formulas and that the word oxide is not preceded immediately by an “a” or an “o”. Worked Example 5.9 Name the following ionic compounds: (a) Fe (SO ) , (b) Al(OH) , and (c) Hg O. Strategy 2 4 3 3 2 Begin by identifying the cation and anion in each compound, and then combine the names for each, eliminating the word ion. Think About It Be careful not to confuse the subscript in the formula with the charge in the metal ion. In part (a), for example, the subscript on Fe is 2, but this is an iron(III) compound. 36 Worked Example 5.10 Name the following species: (a) BrO4, (b) HCO 3 and (c) H C2 . 3 Strategy Each species is either an oxoanion or an oxoacid. Identify the “reference oxidation” (the one with the –ate ending) for each, and apply the rules to determine appropriate names. Think About It Make sure that the charges sum to zero in each compound formula. In part (a), for example, Hg2+ + 2Cl = (+2) + 2(-1) = 0; in part (b), (+2) + 2(-1) = 0; and in part (c), 3(+1) + (-3) = 0. Worked Example 5.11 Determine the formula of sulfurous acid. Strategy The –ous ending in the name of an acid indicates that the acid is derived from an oxoanion ending in –ite. The oxoanion must be sulfite, SO32-, so add enough hydrogen ions to make a neutral formula. Worked Example 5.12 Calculate the molecular mass or the formula mass, as appropriate, for each of the following corresponds: (a) propane, C H , (b) lithium hydroxide, LiOH, and (c) barium acetate, 3 8 Ba(C H O ) . 2 3 2 2 Strategy Determine the molecular mass (for each molecular compound) or formula mass (for each ionic compound) by summing all the atomic masses. 37 Worked Example 5.13 Lithium carbonate, L2 CO3, was the first “mood-stabilizing” drug approved by the FDA for the treatment of mania and manic-depressive illness, also known as bipolar disorder. Calculate the percent composition by mass of lithium carbonate. Strategy Use Equation 5.1 to determine the percent by mass contributed by each element in the compound. Worked Example 5.14 Determine (a) the number of moles of CO 2n 10.00 g of carbon dioxide and (b) the mass of 0.905 mole of sodium chloride. Strategy Use molar mass to convert from mass to moles and to convert from moles to mass. The molar mass of carbon dioxide (CO 2 is 44.01 g/mol and the molar mass of sodium chloride (NaCl) is 58.44 g/mol. Think About It Always double-check unit cancellations in problems such as these–errors are common when molar mass is used as a conversion factor. Also make sure that the results make sense. In both cases, a mass smaller than the molar mass corresponds to less than a mole of substance. 38 Worked Example 5.15 (a) Determine the number of water molecules and the numbers of H and O atoms in 3.26 g of water. 19 (b) Determine the mass of 7.92×10 carbon dioxide molecules. Strategy Use molar mass and Avogadro’s number to convert from mass to molecules, and vice versa. Use the molecular formula of water to determine the numbers of H and O atoms. 2 23
 (b) 7.92×10 CO molecules × 6.022×10 CO molecules × 1 mol CO 
 = 5.79×10 -3g CO 2 Think About It Again, check the cancellation of units carefully and make sure that the magnitudes of your results are reasonable. Worked Example 6.1 Classify the following bonds as nonpolar, polar, or ionic: (a) the bond in ClF, (b) the bond in CsBr, and (c) the carbon-carbon double bond in C H .2 4 Strategy Electronegativity values are: Cl (3.0), F (4.0), Cs (0.7), Br (2.8), C (2.5). Use this information to determine which bonds have identical, similar, and widely different electronegativities. Solution (a) The difference between the electronegativies of F and Cl is 4.0 – 3.0 = 1.0, making the bond in ClF polar. (b) In CsBr, the difference is 2.8 – 0.7 = 2.1, making the bond ionic. (c) In C2H4, the two atoms are identical. (Not only are they the same element, but each C atom is bonded to two H atoms.) The carbon-carbon double bond is C H is 2o4polar. 39 Worked Example 6.2 Burns caused by hydrofluoric acid [HF(aq)] are unlike any other acid burns and present unique medical complications. HF solutions typically penetrate the skin and damage internal tissues, including bone, often with minimal surface damage. Less concentrated solutions actually can cause greater injury than more concentrated ones by penetrating more deeply before causing injury, thus delaying the onset of symptoms and preventing timely treatment. Determine the magnitude of the partial positive and partial negative charges in the HF molecule. Strategy Solve for Q. Convert the resulting charge in coulombs to units of electronic charge. According to Table 6.2, μ = 1.82 D and r = 0.92 Å for HF. The dipole moment must be converted from debye to C·m and the distance between ions must be converted to meters. Worked Example 6.4 Draw the Lewis structure for carbon disulfide (CS ). 2 Setup Step 1: C and S have identical electronegativities. We will draw the skeletal structure with the unique atom, C, at the center. Step


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