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Chemistry I Notes Weeks 1-5

by: Alisha Henderson

Chemistry I Notes Weeks 1-5 CHE 111

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Alisha Henderson


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Elements, Atoms, Molecules, Compounds, Mixtures, Properties of Matter, Types of Changes, Chemical Reactions, Separating Mixtures, Significant Figures, SI Units, Mass, Volume, Length, Temperature, D...
Chemistry I
Dr. Paul
Class Notes
Chemistry, Math
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This 18 page Class Notes was uploaded by Alisha Henderson on Friday September 30, 2016. The Class Notes belongs to CHE 111 at Gordon College taught by Dr. Paul in Fall 2016. Since its upload, it has received 6 views. For similar materials see Chemistry I in Chemistry at Gordon College.


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Date Created: 09/30/16
Chemistry I CHE 111 Instructor(s): ​Dr. Michael Paul ➢ Chemistry is the study of properties and behaviors of matter ➢ It is central to our fundamental understanding of many science related fields ➢ Matter is anything that ha mass and takes up space ○ Elements ○ Molecules ○ Compound (at least 2 atoms) ○ Mixture of elements and a compound ➢ Atoms are the building block of matter ➢ Each element is a unique kind of atom ➢ A compound is made of two or more different kinds of elements ➢ 92 Naturally occurring elements Methods Of Classification ➢ States of matter: solid, liquid, gas ➢ Composition of matter: what is it made of? ➢ Freezing- 0C, 32F, 273K ➢ Classification based on composition ○ Homogeneous Mixture (same throughout) ○ Heterogeneous Mixture ○ Element ○ Compound Classification of Matter-Substance ➢ A substance has distinct properties and a composition that does not vary from sample to sample ➢ The two types of substances are elements and compounds ➢ Element: Substance that can NOT be decomposed ➢ Compound: Substance that CAN be decomposed Compounds and Composition ➢ Compounds have a definite composition that means the relative number of atoms of each element that makes up the compound is the same in any sample ➢ THis is the Law Of Constant Composition Classification of Matter- Mixture ➢ Mixtures exhibit properties of the substances that make them up ➢ Mixtures can vary in composition throughout a sample (hetero) or have the same composition throughout the sample (homo) ➢ Another name for homo mixture is a solution Monday, August 29, 2016 Section 1.3-Properties of Matter ➢ Physical Properties​ can be observed without changing a substance into another substance. ○ Some examples include boiling point, density, mass, or volume ➢ Chemical Properties​ can ONLY be observed when a substance is changing into another substance ○ Some examples include flammability, corrosiveness, or reactivity with acid ➢ Intensive Properties​ are independent of the amount of the substance that is present ○ Examples include density, boiling point or color ➢ Extensive Properties​ depend upon the amount of the substance present ○ Examples include mass, volume, or energy Types of Changes ➢ Physical Changes​ are changes in matter that do NOT change the composition of a substance ○ Examples include changes of state, temperature, and volume ➢ Chemical Changes​ result in new substances ○ Examples include combustion, oxidation, and decomposition Changes in State of Matter ➢ Converting between the 3 states is physical change ➢ When ice melts or water evaporates there are still 2 H atoms and 1 O atom in each molecule Chemical Reactions (Chemical Change) ➢ In the course of a chemical reaction, the reacting substances are converted to new substances. Separating Mixtures ➢ Mixtures can be separated based on physical properties of the components of the mixture ○ Examples: filtration, distillation, chrom ➢ Filtration​: solid substances are separated from liquids and solutions ○ “Coffee filter” ➢ Distillation​: uses differences in the boiling points of substances a separate a homogeneous mixture into its components ○ Boiling salty water. Water will boil first and turn into a gas ➢ Chromatography: ​ The technique separates substances on the basis of differences in the ability of substances to adhere to the solid surface, in this case, dyes the paper. Different solubilities. Section 1.4 Numbers and Chemistry ➢ Numbers play a major role in chemistry. Many topics are quantitative (How much) ➢ Concepts of numbers in science ○ Units of measurement ○ Quantities that are measured and calculated ○ Uncertainty in measurement ○ Significant figures ○ Dimensional analysis SI Units ➢ The International System of Units ➢ Base Units ○ Kg-mass ○ M-length ○ S or sec-time ○ K-temp ○ Mol-amount of substance ○ A or Amp-electric current ○ Cd-luminous intensity ➢ Metric System ○ G-mass ○ M-length ○ S or sec-times ○ C or K-temperature ○ Mole-amount of a substance ○ Cc or cm^3- volume ➢ Prefixes: ○ Kilo-1,000 ○ Decu- 1/10 ○ Centi 1/100 ○ Deci- 10 Mass and Length ➢ These are basic units ➢ Mass: measure of the amount of material in an object. Kilogram(SI) and gram(Metric) ➢ Length: distance. Meter Volume ➢ Volume is not a base unit for SI. it is derived from length (m x m x m =m^3) ➢ The most commonly used metric units for volume are the liter (L) ➢ A liter is a cube 1 decimeter (dm) long on each side ➢ Milliliter is a cube 1 centimeter (cm) long on each side, also called 1 cubic centimeter (cm x cm x cm =cm^3) Wednesday, August 31st Volume Conversions ➢ 1 decimeter = 0.1 meter ➢ 10 dm = 1 m ➢ 100 dm^2 = 1 m^2 ➢ 1,000 dm^3 = ​1,000 Liter = 1 m^3 ➢ 1,000 milliliters = 1 Liter ➢ 100 cm = 1 m ➢ 10,000 cm^3 = 1 m^2 ➢ 1,000,000 cm^3 = 1 m^3 = 1,000 Liters ➢ 1,000 cm^3 = 1 Liter Temperature ➢ In general usage, temperature is considered the “hotness and coldness” of an object that determines the direction of heat flow ➢ Heat flows spontaneously from an object with a higher temperature to an object with a lower temperature ➢ 373K = 100C = 212F ➢ 310K = 37C + 98.6F ➢ In scientific measurements use C or K ➢ The Celsius scale is based on the properties of water ○ 0C is the freezing point of water ○ 100C is the boiling point of water ➢ K is the SI unit of temperature ○ Based on properties of gases ○ There are no negative K ○ The lowest possible tem f called absolute zero (0K) ➢ K = C + 273.15 ➢ The F scale is not used in scientific measurements, but you hear about it in weather reports ➢ The equations below allow for conversion between F and C ○ F = 9/5(C) =32 ○ C = 5/9(F - 32) Density ➢ Density is a physical property of a substance ➢ It has units that are derived from the units for mass and volume ➢ The most common units are g/mL or g/cm^3 ➢ D=m/V ➢ M = 65 ➢ D = 0.791 ➢ V = m/V Section 1.5 ➢ Exact​ numbers are counted or given by definition. For examples there are 12 eggs in 1 dozen ➢ Inexact ​(or measured) numbers depend on how they are determined. Scientific instruments have limitations. Some balances measure to +/- 0.01g others measure to +/- 0.0001g Uncertainty in Measurements ➢ Different measuring devices have different uses and different degrees of accuracy ➢ All measured numbers have some degree of inaccuracy ○ Burette is most accurate, graduated cylinder is least Accuracy vs Precision ➢ Accuracy​ refers to the proximity of a measurement to the true value of a quantity ➢ Precision​ refers to the proximity of several measurements to each other Significant Figures ➢ The term significant figures refers to the digits that were measured ➢ Rounding numbers ➢ RULES: ○ All nonzero numbers are significant ○ Zeros between two significant figures are themselves significant ○ Zeros at the beginning of a number are never significant ○ Zeros at the end of a number are significant if a decimal point is written in the number ■ 100 = significant figure is 1 ■ 100. = 3 sig fig ■ 40.000 = 5 sig fig ➢ When addition or subtraction is performed, answers are rounded to the least sig decimal place ○ 23.4​ + 0.54 =23.94 =​23.9 ➢ When multiplication or division is performed, answers are rounded to the digits that correspond to the least number of sig figs ○ 3.32 x 2.0 = 6.64 = 6.6 Friday, September 2nd Dimensional Analysis ➢ Uses conversion factors ➢ We use the ratio that allows us to change units ➢ (miles) / (hr) x (meter) / (miles) x (1hr) / (3600sec) ➢ (515m) / (s) x (3600s) / (1hr) x (1 mile) / (1609.344 m) = 1152.02 = 1150 mph Chapter 2: Atoms, Molecules and Ions Atomic Theory of Matter ➢ The theory that atoms are fundamental building blocks of matter reemerged in the early nineteenth century, championed by John Dalton ➢ Dalton’s Atomic Theory (Postulates) ○ Each element is composed of small particles called atoms ○ All atoms of a given element are identical, but the atoms are one element are different from others ○ Atoms of one element cannot be changed into atoms of another ○ Compounds are formed when atoms of more than one element combine Law of Conservation of Mass ➢ The total mass of substances at the end of a process is the same as the mass in the beginning Law of Multiple Proportions ➢ Same ratio Wednesday, September 7th Discovery of Subatomic Particles ➢ In his view the atom was the smallest particle possible. Many discoveries led to the fact the atom itself was made of smaller particles ○ Electrons and cathode rays ○ Radioactivity ○ Nucleus, protons, neutrons The Electron (Cathode Ray) ➢ Streams of negatively charged particles were found to emanate from cathode tubes, causing fluorescence ➢ JJ Thomson​ is credited with their discovery (​1897​) ➢ Electrons move from the negative cathode to positive ➢ The tube contains a fluorescent screen that shows path ➢ The rays are deflected by a magnet ➢ JJ measured the charge/mass ratio of the electron to be 1.76 x 10^8 coulombs/gram (C/g) Millikan Oil-Drop Experiment (Electrons) ➢ Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other ➢ Robert Millikan determined the charge on the electron in 1909 ➢ Charge of Electron: -1.602 x 10^19 Radioactivity ➢ Radioactivity is the spontaneous emission of high-energy radiation by an atom ➢ It was first observed by Henri Becquerel ➢ Marie and Pierre Curie also studied it ➢ Its discovery showed that the atom had more subatomic particles and energy associated with it ➢ Three types of radiation were discovered by Ernest Rutherford: ○ A particles (positively charged) ○ B particles (negatively charged, like electrons) ○ Y rays (uncharged) The Atom, circa 1900 ➢ The prevailing theory was that of the “plum pudding” model, put forward by Thomson ➢ It featured a positive sphere of matter with negative electrons embedded in it Discovery of the Nucleus ➢ Ernest Rutherford shot a particles at a thin sheet of gold foil and observed the pattern of scatter of the particles The Nuclear Atom ➢ Since some particles were deflected at large angles, Thomson’s model could not be correct ➢ Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom ➢ Most of the volume is empty space ➢ Atoms are very small ○ 1 - 5 A or 100 - 500 pm ➢ Other subatomic particles (protons and neutrons) were discovered Subatomic Particles ➢ Protons (+1) and electrons (-1) have a charge; neutrons are neutral ➢ Protons and Neutrons have essentially the same mass (relative mass 1). The mass of an electron is so small we ignore it (Mass 0) Symbols of Elements ➢ Atomic number = number of protons. Subscript, Z ➢ Mass Number = total protons and neutrons. Superscript, A Friday, September 9th Atomic Mass ➢ Atoms have extremely small masses ➢ The heaviest known atoms have a mass of 4 x 10^-22 g ➢ A mass scale on the atomic level is used, where an atomic mass unit (amu) is the base unit ➢ 1 amu = 1.66054 x 10^-24 g Atomic Weight Measurement ➢ Atomic and molecular weight can be measured with great accuracy using a mass spectrometer ➢ Masses of atoms are compared to the carbon atom with 6 protons and 6 neutrons (C-12) Isotopes ➢ Isotopes ​are atoms of the same element with different masses ➢ Isotopes have different numbers of neutrons, but same number of protons Atomic Weight ➢ Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations ➢ An average mass is found using all isotopes of an element weighted by their relative abundances. Periodic Table ➢ Elements are ordered by atomic number ➢ Atomic number is on the bottom (protons) ➢ Atomic weight is on the top (protons and neutrons) ➢ Columns are groups (1-18) ➢ Periods are horizontal ➢ Groups: ○ 1A ​= alkali metals = Li, Na, K, Rb, Cs, Fr ○ 2A ​= Alkaline earth metals = Be, Mg, Ca, Sr, Ba, Ra ○ 7A​ = Halogens =F, Cl, Br, I, At ○ 8A ​= Noble Gases (rare gases) = He, Ne, Ar, Kr ➢ Periodicity​- when one looks at the chemical properties of elements, one notices a repeating pattern of reactivities ➢ Metals are on the left ➢ Metal properties include: ○ Shiny luster ○ Conducting heat and electricity ○ Solidity (except mercury) Monday, September 12th Chemical Formulas ➢ The subscript to the right of the symbol of an element tells the number of atoms of that element is one molecule of the compounds ➢ Molecular compounds​ are composed of molecules and almost always contain only nonmetals Diatomic Molecules ➢ These seven elements occur naturally as molecules containing two atoms: ○ Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, Iodine Types of Formulas ➢ Empirical formulas​ give the lowest whole-number ratio of each element in a compound ○ CH ➢ Molecular formulas​ give the exact number of atoms of each element in a compound ○ C6H6 ➢ If we know the molecular formula of a compound, we can determine its empirical formula. The converse is not true! ➢ Structural formulas​ show the order in which atoms are attached. They do NOT depict the 3-D shape of molecules ➢ Perspective drawings​ also show the 3-D order of the atoms in a compound. These are also demonstrated using ​models. 2.7: Ions and ionic Compounds Ions ➢ When an atom of a group of atoms loses or gains electrons, it becomes an ion ➢ Cations are when an electron is lost. Monatomic cations are formed by metals ➢ Anions are when an electron is gained. Monatomic anions are formed by non metals ➢ KNOW COMMON CATION AND ANIONS TABLE!!!!! ➢ 2+ is the most common stable state Ionic Compounds ➢ Ionic compounds (such as NaCl) are generally formed between metals and nonmetals ➢ Electrons are transferred from the metal to the nonmetal. The oppositely charged ions attract each other. Only empirical formulas are written. Writing Formulas ➢ Because compounds are electrically neutral, one can determine the formula of a compound this way: ○ The charge on the cation becomes the subscript on the anion ○ The charge on the anion becomes the subscript on the cation ○ If these subscripts are not the lowest whole number ratio, divide by the GCF ○ Metal and a nonmetal = switch the superscripts Wednesday, September 14th Inorganic Nomenclature ➢ Write the name of the cation. If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses. ➢ If the anion is an element, change its ending to ​-ide​. If the anion is a polyatomic ion, simply write the name of the polyatomic ion. ○ CuO = copper (II) oxide ○ MgO = magnesium oxide ○ FeO = Iron (II) Oxide ○ Fe2O3 = Iron (III) Oxide ○ FeS = Iron (II) Sulfide ○ NaCl = Sodium Chloride Oxyanion Nomenclature ➢ When there are two oxyanions involving the same element ○ The one with fewer oxygens ends in ​-ite ○ The one with more oxygens ends in ​-ate ➢ Central atoms on the second row have a bond to, at most, three oxygens; those on the third row take up to four ➢ Charges increase as you go from right to left Friday, September 16th Acid Nomenclature ➢ If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- ○ HCl: hydrochloric acid ○ HBr: hydrobromic acid ○ HI: Hydroiodic acid ➢ If the anion ends in -ite, change the ending to -ous acid ○ HCLO: hypochlorous acid ○ HCLO2: Chlorous acid ➢ IF the anion ends in -ate, change the ending to -ic acid ○ HCLO3: chloric acid ○ HCLO4: perchloric acid ➢ Examples: HCN = H+ and CN- = Cyanide =hydrocyanic acid ➢ Examples: HNO3 = H+ and NO3- =nitric acid Nomenclature of Binary molecular compounds (nonmetal and a nonmetal) ➢ The name of the element father to the left in the periodic table (closer to the metals) or lower in the same group is usually written first ➢ A prefix is used to denote the number of atoms of each element in the compound (Mono- is not used on the first element) ➢ The ending f the second element is changed to -ide ○ CO2: carbon dioxide ○ CCL4: carbon tetrachloride ➢ IF the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one ○ N2O5: 3.1 Chemical Equations Stoichiometry ➢ The study of the mass relationship in chemistry ➢ Based on the Law of Conservation of Mass (Antoine Lavoisier, 1789) ➢ Chemical equations are concise representations of chemical reactions ➢ Reactant → Products ➢ The states of the reactants and products are written in parentheses to the right ➢ Coefficients are inserted to balance the equation to follow the law of conservation of mass Monday, September 19th, 2016 Combination reactions ➢ In combination reactions two or more substances react to form one product ➢ In decomposition reaction one substance breaks down into two or more substances ➢ Combustion Reactions are generally rapid reactions that produce a flame ○ Most often involve oxygen in the air as a reactant Wednesday, September 21st Avogadro’s number ➢ In a lab, we cannot work with individual molecules. They are too small. ➢ 6.02 x 10^23 ​atoms or molecules is an amount that brings us to lab size. It is ONE MOLE ➢ One mole of 12C has a mass of 12.000g Molar Mass ➢ A molar mass is the mass of 1 mol of a substance ➢ The molar mass of an element is the atomic weight for the element from the periodic table. If it is diatomic, it is twice that atomic weight. ➢ The formula weight (in amu) will be the same number as the molar mass (in g/mol) Using Moles Grams​ ←→ use molar mass ←→ M ​ oles​←→ avogadro's # ←→ f ​ ormulas units Determining Empirical Formula Mass % of elements → ​ assume 100g → ​Grams of each Element​ → use molar mass → ​Moles of each Element ​→ calculate mole ratio → e ​ mpirical formula Determining Molecular formula ➢ The number of atoms in a molecular formula is a multiple of the empirical ➢ If we find the empirical formula and know a molar mass (Molecular weight) for the compound, we can find the molecular formula ➢ Whole-number multiple = (molecular weight) / (empirical formula weight) Combustion Analysis ➢ Compounds containing C, H, and O are routinely analyzed through combustion in a chamber Monday, September 26th Quantitative Relationships ➢ The coefficients in the balanced equation show ○ Relative numbers of molecules of reactants and products ○ Relative numbers of moles of reactants and products, which can be converted to mass Stoichiometric Calculations Grams of substances A→ ​ use molar mass of A→ ​Moles of Substance A​→ use coefficients from balanced equation→ ​Moles of Substance B​→ use molar mass of B→ Grams of Substance B ➢ Example: ​ How many grams of water can be produced from 1.00g of glucose? ○ C6H12O6(s) + 5O2(g) → 6CO2 (g) + 6H2O(l) 3.7 Limiting Reagents Limiting Reactants ➢ The limiting reactant is the reactant present in the smallest stoichiometric amount ○ In other words, it’s the reactant you'll run out of first ➢ Examples: ○ N2 + 3H2 → 2NH3, How many moles of NH3 can be formed from 3.0 mol of N2 and 6 Mol of H2? ○ Limiting Reactant is H2 ○ Moles H2 = (3 mol N2) x (3 mol H2)/(1 mol N2) = 9 mol H2 Theoretical Yield ➢ The theoretical yield is the maximum amount of product that can be made ○ In other yields, its the amount of product possible as calculated through the stoichiometry problem ➢ This is different than actual yield Percent Yield ➢ One finds the percent yield by comparing the amount actually to the amount possible ➢ (Actual yield) / (theoretical yield) x 100 = percent yield Wednesday, September 28th, 2016 Solutions ➢ Solutions​ are defined as homogeneous mixtures of two or more pure substances ➢ The ​solvent​ is present in greatest abundance ➢ All other substances are ​solutes ➢ When water is the solvent, the solution is called an ​aqueous solution Aqueous Solutions ➢ Substances can dissolve in water by different ways ○ Ionic Compounds dissolve by dissociation, where water surrounds the separated ions ○ Molecular compounds interact with water, but most do NOT dissociate ○ Some molecular substances react with water when they dissolve Electrolytes and Nonelectrolytes ➢ An electrolyte is a substance that dissociates into ions when dissolved in water ➢ A nonelectrolytes may dissolve in water, but it does not dissociate into ions when it does Strong Electrolyte Weak Electrolyte Nonelectrolyte Ionic All None None Molecular Strong acid Weak acid All other compound Electrolytes ➢ A strong electrolyte dissociates completely when dissolved in water ➢ A weak electrolyte only dissociates partially when dissolved in water ➢ A nonelectrolyte does NOT dissociate in water, does not conduct electricity ➢ Electrolyte does conduct electricity 4.2 Precipitation Reactions Precipitation Reactions ➢ When two solutions containing soluble salts are mixed, sometimes an insoluble salt will be produced. A salt “falls” out of solution, like snow out of the sky. This solid is called a ​precipitate ➢ Look for (s). A Solid. Solubility of Ionic Compounds ➢ Not all ionic compounds dissolve in water ➢ A list of ​solubility rules​ is used to decide what combination of ions will dissolve ➢ The solubility rules will be given to us on the test. Be able to use. Metathesis (Exchange) Reactions ➢ Metathesis comes from the Greek word that means “to transpose” ➢ It appears as though the ions in the reactant compounds exchange, or transpose, ions ➢ Ag​NO3 (aq) + K​Cl ​(aq) → ​AgCl​ (s)​ KNO3 (aq) Completing and Balancing Metathesis Equations 1. Use the chemical formulas of the reactants to determine which ions are present 2. Write formulas for the products: cation from one reactant, anion from the other. Use charges to write proper subscripts 3. Check solubility rules. If either product is insoluble, a precipitate forms 4. Balance the solution Friday, September 30th Molecular Equation ➢ Molecular Equation lists the reactants and products without indicating the ionic nature of the compounds Complete Ionic Equation ➢ In the ​complete ionic equation​ all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions ➢ This more accurately reflects the species that are found in the reaction mixture ➢ Ag+(aq) + No3-(aq) + K+(aq) + Cl-(aq) → AgCl(s) + K+(aq) +No3-(aq) Net Ionic Equation ➢ To from the net ionic equation, cross out anything that does not change from the left side of the equations to the right ➢ The ions crossed out are called ​spectator ions​, K+ and NO3-, in this example ➢ The remaining ions are the reactants that form the product-- an insoluble salt in a precipitation reaction as in this example ➢ Ag+(aq) + Cl-(aq) → ​AgCl(s) Writing Net Ionic Equations 1. Write a balanced molecular equations 2. Dissociate all strong electrolytes 3. Cross out anything that remains unchanged from the left side to the right side 4. Write the net ionic equation with the species that remain Example​: Calcium chloride and sodium carbonate are mixed. Write a net ionic equation 1. CaCl2(aq) + Na2CO3(aq) → ​CaCO3(s)​ + 2NaCl(aq) 2. Ca2+(aq) + 2Cl-(aq) + 2Na+(aq) + CO3^2-(aq)→CaCO3(s)+2Na+(aq) + 2Cl-(aq) 3. Cross out 2Cl- and 2Na+ 4. Ca+(aq) + CO3^2- → CaCo3(s) (NET IONIC) 4.3 Acids, Bases and Neutralization Acids ➢ The Swedish physicist and chemist S.A. Arrhenius defined acids as substances that increase the concentration of H+ when dissolved in water ➢ Both the Danish chemist J.N. Bronsted and the British chemist T.M. Lowry defined them as proton donors Bases ➢ Arrhenius defined bases as substances that increase the concentration of OH- when dissolved in water ➢ Bronsted and Lowry defined them as proton acceptors Strong or Weak Acids ➢ Strong acids completely dissociate in water; weak acids only partially dissociate ➢ Strong bases dissociate to metal cations and hydroxide anions in water; weak bases only partially react to produce hydroxide anions ➢ CHART 4.2 Common Strong Acids and Bases Acid-Base Reactions ➢ In an acid-base reaction, the acid (H20) donates a proton (H+) to the base (NH3) ➢ Reactions between an acid and base are called neutralization reactions ➢ When the base is a metal hydroxide, water and a salt (an ionic compound) are produced Neutralization Reactions ➢ When a strong acid (like HCl) reacts with a strong base (like NaOH), the net ionic equation is circled below: ○ HCl(aq) + NaOH(aq) → NaCl(aq) + H2O ○ H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → Na+(aq) +Cl-(aq) +H20(l) (cancel similar) ○ H+(aq) + OH-(aq) → H2O Gas-Forming Reactions ➢ Some metathesis reactions do not give the product expected ➢ When carbonate or bicarbonate reacts with an acid, the products are a salt, carbon dioxide, and water ➢ CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l) ➢ NaHCO3(aq) + HBr(aq) → NaBr(aq) +Co2(g) + H2O(l)


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