Chemistry Minimum Learning Outcomes – The student must be able to demonstrate proficiency in:

Basic Units of measurements:1. Length : the unit is meter (m)2. Volume : the units are Liter (L), milliliter (mL), cm3 or cc which is equal to a milliliter3. Mass : the unit is gram (g) or kilogram (kg)4. Temperature : units are Celsius (C) and Kelvin (K)5. Energy : Joule (J)6. Amount of a substance : mole

Scientific Notation: is writing numbers in an exponential form (the power of ten) 1234 = 1.234 x 103

Prefixes and Equalities: know the prefixes of the Metric and SI (International System) Example: 1 kilometer (km) = 1000 meters (m). kilo is the prefix which is equal to 1000.

Conversion factors: are qualities written in the form of a fraction, where one of the quantities is the numerator, and the other is the denominator.Example: 12 in = 1 ft, this equality is written as a conversion factor such as: 12 in or 1 ft 1ft 12 in

Problem solving using conversion factorsQuantity (given unit) x conversion factor (desired unit/given unit) = desired unit

Example 1: How many moles are in 45 grams of sodium?Plan: g à mole45 g Na x 1 mole Na = 1.96 mole Na 22.99 g NaExample 2: A chemist needs 325 mg of potassium. If the only source of potassium available was 0.50 g/mL of solution, how many mL of solution would she need?Plan: mg à g à mL325 mg Potassium x 1 g x 1 mL solution = 0.65 mL solution 1000 mg 0.50 g Potassium

Density: is the mass per unit volume.Density (D) = Mass(M)/volume(V), or [D=M/V]*Density is a physical property that can be used to help identify an unknown substance.*Difference in density determines whether an object will sink or float. The denser object will sink and

the less dense object will float.Units of density: density of solid g/cm3, density of liquids g/mL, density of gases g/L.

Example 1: What is the density of a piece of wood that weighs 3.00 g and has a volume of 5.55 cm3.D = M/V = 3.00 g/ 5.55 cm3 = 0.541 g/cm

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Example 2: What is the mass (g) of a liquid that has a density of 0.955 g/mL and a volume of 15.0 mL. D=M/V so M= DxV= 0.955 g/mL x 15.0 mL = 14.3 gExample3: Calculate the volume of a gas that has a mass of 2.50 g and a density of 0.587 g/L.D = M/V so V = M/D= 2.50 g / 0.587 g/L = 4.26 L

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TemperatureTemperature: is the measure of heat in an object.There are three different scales to measure temperature, and each one has its own units.

1. Celsius scale: Degree C2. Kelvin: K

The relationship between the three scales of temperature:K= C + 273 or C = K – 273Examples: Convert 33.0 C to K. answer: K = C + 273 = 33.0 + 273 = 306 KExample: Convert 450. K to C. answer: C= K – 273 = 450. – 273 = 177 C

Elements and symbolsElements: are primary substances from which all other substances are built and they cannot be broken

down into simpler substances by ordinary means.Chemical symbols: are one or two letter abbreviations for the names of the elements.

The Periodic Table: is made of groups (vertical lines) and periods (horizontal lines). Groups: there are 18 groups which are classified to two parts:

1. Representative elements: are called the “A” groups, and numbered with Roman numerals as IA, IIA, IIIA,…or as 1A, 2A, 3A…

2. Transition elements: are called “B” groups and numbered with Roman numerals as IB, IIB, IIIB,.. or as 1B, 2B, 3B,…

Naming of some groups: 1) Alkali metals: are the elements in group 1A2) Alkaline earth metals: are the elements in group 2A3) Halogens: are the elements in group 7A4) Noble gases: are elements in group 8A.5) Transition elements: are all elements in groups B.

Elements are also classified according to their physical properties to:1. Metals – The elements underneath the bold “stair-steps”2. Nonmetals – The elements above the bold “stair-steps”3. Metalloids – The elements along the bold “stair-steps” that have properties of both metals

and non-metals.

Dalton Atomic theory:-All matter is made up of tiny particles called atoms.-All atoms of a given element are similar to one another.-Atoms of two or more different elements combine to form compounds.

J.J. Thompson’s Atomic Model:-Atoms have negative subatomic particles called electrons.

Rutherford’s Atomic Model:-Atoms have their positive particles in a central condensed area called a nucleus.

Bohr’s Atomic Model:-Atoms have located electrons in definite orbits around a central nucleus.

Structure of the atom: it has nucleus where protons and neutrons reside. The electrons are located in clouds around the nucleus in different energy levels.

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The Atom: is the smallest particle of an element that retains the characteristics of that element. It consists of subatomic particles:

Location Charge Mass, amuproton in nucleus +1 1neutron in nucleus 0 1electron in orbit around nucleus -1 1/2000

Atomic number (Z): the number of protons in the nucleus.In the neutral atom the number of electrons equals the number of protons.

Mass number (A): the sum of the number of protons and neutrons in the nucleus.Isotopes: atoms of the same element that have different numbers of neutrons.

Atomic symbols for isotopes of Hydrogen: 11H, (top number is the mass number-A, and lower number is

the atomic number-Z), 21H, and 3

1H.Example: identify the number of protons and neutrons in the 25

12Mg isotope.Answer: 12 protons and 25-12= 13 neutrons.

Atomic mass (weight): is the average mass of all of the naturally occurring isotopes of an element measured in atomic mass units (amu). They are written underneath the symbol of the element in the periodic tableExample: atomic mass of H is 1.008 amu, for C is 12.01 amu.

Electron energy levels: electrons spin around the nucleus in a certain space called energy levels (or shells). Energy levels are labeled (n). The first energy level that is closer to the nucleus has n=1 (holds a maximum of 2 electrons) and the next has n=2, (holds a maximum of 8 electrons….Maximum number of electrons in an energy level is 2n2.Energy levels are subdivided to orbitals, where each orbital can hold a maximum of two electrons. There are four different kinds of orbitals, which are classified according to their shapes: s, p, d, f

Electron shell configuration: you need to be able to write these for elements 1 - 56.Example: what is the electron configuration of fluorine? What is the noble gas electron configuration?Answer: 1s2 2s2 2p5 noble gas electron configuration: [He] 2s2 2p5

The Periodic Law: is the regular pattern of change in physical and chemical properties of elements as their atomic number increases.

Compounds and Their BondsValance electrons: Are electrons located in the valence shell, which is the outermost energy level of an

atom.

Electron-Dot structure: is the symbol of the element with valence electrons shown as dots. Like H·, He:, Li·,.. (Each dot represents one valence electron)

Octet Rule: atoms tend to lose, gain, or share electrons to achieve the valence electron configuration of the noble gases (8e- except those elements reaching the electron configuration of helium which holds 2e-).

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Ions: are atoms or group of atoms that have lost or gained electrons, and are classified according to their charges:Cations (Positive ions): are atoms that have lost electron(s). Examples: Li loses 1e- to become Li+ ion, Ca loses 2e- to become Ca+2 ion, Al loses 3e- to become Al+3 ion,…Anions (Negative ions): are atoms that have gained electron(s). Examples: F gains 1e- to become F- ion, O gains 2 e- to become O-2 ion, N gains 3 e- to become N-3 ion.,..

Ionic charges from group number:Examples: write the ionic forms for Na (Group 1A), Mg (Group 2A), Cl (Group 7A).Answer: Na+, Mg+2, Cl-.

Ionic CompoundsIonic compounds: are formed between Metals and Non-metals when the valence electrons are

transferred. (Metal) cations and (nonmetal) anions are held together by an electrostatic attraction between opposite charges in what is called an ionic bond. Example: Li+ is attracted to F- to form LiF compound.

The chemical formula: it indicates the number and kinds of atoms that make up the compound. Like; LiF, NaCl, CaF2,..

Charge balance in an ionic compound: the negative and positive charges have to be the same. Na+ is attracted to F- and the formula is written as NaF. 2 Na+ are attracted to O-2 and the formula is written as Na2O and the 2 is called “subscript”.Hint: “Crisscross” the oxidation numbers (without the + or -) to get the subscripts.Examples: write the formula that will result of the reaction between: a) Oxygen and Lithium. b) Sulfur and calcium.Answer: a) Li2O b) CaS

Naming ionic compounds:1) When metal ions have fixed charges (1A Metals): CaBr2 is called calcium bromide. Al2S3 is

called aluminum sulfide. Ca3N2 is called calcium nitride.2) When metal ions have variable charges (transition metals and a couple others): FeF2 is called

iron (II) fluoride, while FeF3 is called iron (III) fluoride. Cu2O is called copper (I) oxide, while CuO is called copper (II) oxide.Examples: write the chemical formulas for the following compounds: a) Lead (II) chloride. b) Sodium bromide. c) Potassium Nitride.

Answer: a)PbCl2 b) NaBr c) K3N

Polyatomic Ions: are group of atoms that have an electrical chargeExamples: OH- (hydroxide), NH4

+ (ammonium),….see Reference Table

Writing formulas for compounds containing polyatomic ions and naming:Examples: CaCO3 (calcium carbonate), K2NO3 (potassium nitrate), Ag2SO4 (silver sulfate),…

Writing a formula from a name:Examples: sodium hydrogen carbonate aluminum sulfate magnesium phosphate First write the ions: Na+ HCO3

- Al+3 SO4-2 Mg+2 PO4

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Balance charges (crisscross): NaHCO3 AL2(SO4)3 Mg3(PO4)2

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Covalent CompoundsCovalent Compounds: are formed when a nonmetal combines with another nonmetal by sharing some

of their electrons. The resulting bond is called the covalent bond. Like; H combines with F to form HF compound, H combine with H to form H2 molecule, N combine with 3H to form NH3 compound,…

Multiple covalent bonds: covalent bonds are classified to;Single bond: when the two atoms share one pair of electrons, H : HDouble bond: when the two atoms share 2 pairs of electrons, O::O, or C::CTriple bond: when the two atoms share 3 pairs of electrons, N:::N, or C:::O

Naming covalent compounds: subscripts indicating two or more atoms of an element are expressed as prefixes placed in the front of each name .

Examples: CO carbon monoxide, CO2 carbon dioxide, NO nitrogen monoxide, N2O2 dinitrogen dioxide,…

Writing the formulas from names: Examples: sulfur trioxide carbon tetrachloride silicon dioxide SO3 CCl4 SiO2

Electronegativity: is the ability of an atom to attract electrons in a molecule.

Bond Polarity: Covalent bonds also can be classified as:1. Nonpolar covalent bond: formed when atoms with the same or similar electronegativity share

electrons. Like; H-H, H-C, Br2

2. Polar covalent bond: formed when atoms with different electronegativities share electrons. Like H-F, N-O, C-Cl

Examples: predict the type of bonding (ionic, polar covalent, or nonpolar covalent): NH3 K2O Br2 HCl Answer: polar covalent ionic nonpolar covalent polar covalent

Shapes and polarity of moleculesTo determine the shape of a molecule we use the valence-shell electron-pair repulsion theory

(VSEPR). Examples: What are the shapes of the following molecules? CH4 NH3 H2O BH3 CBr4

Answer: tetrahedral trigonal pyramidal bent trigonal planar tetrahedral

Polarity of molecules: molecules are classified to:1) Nonpolar molecule: When all bonds are nonpolar or if the polar bonds have symmetrical

arrangements. Ex: CH4, CO2

2) Polar molecule: When the molecule has only one polar bond or when polar bonds do not cancel each other. Ex: HCl, H2O.

Chemical reactions and quantitiesPhysical change: is when a physical property such as the appearance of a substance is altered, but not

its composition (chemical makeup).Chemical change: when substances change to a new different substances by a chemical reaction.Chemical reaction: it involves a chemical change.

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Chemical equations: show the materials that are used (reactants) and the materials that are formed (products).Reactants à Products

Balancing chemical equations: the chemical equations are balanced by using coefficients in front of some of the formulas… NOT by changing subscripts.

CH4 + 2 O2 à CO2 + 2 H2OThe underlined numbers are used to balance the number of atoms in both sides of the equation.

Types of Reactions: (See Reference Table)1) Synthesis reaction : two or more elements or simple compounds bond together to form product,

example: N2 + 2O2 à 2NO2

2) Decomposition reaction : a reactant splits into tow or more simple products, example:CaCO3 à CaO + CO2

3) Single replacement reaction : Na + HCl à NaCl + H2

4) Double replacement reaction : BaCl2 + Na2SO4 à BaSO4 + 2NaCl5) Combustion reaction : fuel and oxygen react to produce carbon dioxide and water.

CH4 + 2O2 à CO2 + H2O

Oxidation-Reduction Reactions (LEO GER)Oxidation: is loss of electrons. Na à Na+ + e-Reduction: is gain of electrons. O + 2e- à O2-

Example: Ca + S à CaS in this reaction “Ca” is oxidized and “S” is reduced.

The MoleThe mole is a unit used to count the amount of a substance.

For example: one mole of carbon atoms contains 6.02 x 1023 atoms of Carbon.One mole of hydrogen atoms in 1 mole hydrogen gas molecules (H2) is 2 x 6.02 x 1023

1 Mole = mass of element in grams / molecular weight or atomic weight.1 mole of Carbon (C) is 12.01 g.1 mole of hydrogen gas (H2) is 2 * 1.008 g

Avogadro’s number: is equal to 6.02 x 1023 particles for 1 mole of anything.

Moles of elements in a formula:Al2(SO4)3 contains 2 moles Al, 3 moles S, and 12 moles O

Molar Mass: is the mass of one mole of the element in grams, which equals to the atomic weight of the element.Examples: molar mass of carbon is 12.01, molar mass of H2 is 2(1.008) = 2.016, molar mass of SO3 is (32.07) + 3(16.0) = 80.07 g.

Calculations using molar mass: molar mass is used to convert between moles and grams.Example: How many grams are there in 0.780 moles of silver?Answer: from the periodic table use the conversion factor (1 mole Ag = 107.9 g)0.750 moles Ag x 107.9 g Ag = 80.9 g Ag 1 mole Ag) Example: How many moles are there in 3.50 grams KI?Answer: from the periodic table use the conversion factor (1 mole KI = 166.0 g KI)3.50 g KI x 1 mole KI = 0.0211 moles KI

166.0 g KI

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Mole relationships in chemical equationsFrom the balanced chemical equation we can calculate the amount of product or reactant in grams or in

moles using the following diagram:Grams A ↔ Moles A ↔ Moles B ↔ Grams B

Energy and MatterEnergy can be classified as:

Potential energy: stored energy such as chemical energy.Kinetic energy: energy of motion, like walking, or thermal energy…Heat: is the energy associated with the motion (Kinetic Energy) of particles in a substance. Units of heat joule (J)

Specific heat: is the amount of heat that raises the temperature of 1 g of a substance by 1 ºC. The units are J/g.ºC.

Calculations using specific heat: Q = mCp∆T heat = (mass) x (specific heat) x (temperature change, ∆T)

Example: Using your reference tables, identify the probable composition of a material that requires 3588 Joules of heat to raise the temperature of 50 grams of the substance from 20 ºC to 100 ºC.

Cp = Q m x ∆T

= 3588 J 50g x 80 ºC

= 0.897 J/gºCThe material is most likely made of Aluminum since has a specific heat of 0.897 J/gºC

Calorimeter: is a device used to measure the energy transfer between objects.

States of Matter: Solid, Liquid, GasChanges of states:

1) Melting point : is the temperature at which a solid will melt.2) Freezing point : is the temperature at which a liquid will solidify

Sublimation: is the endothermic change of solid to a gas without going through the liquid state.Evaporation: is the endothermic change of liquids to gases.Boiling point: is the temperature at which liquids become a gas (vapor).Condensation: is the exothermic change of vapor to a liquid.

Heat of fusion: is the heat needed to melt a solid.

Heat required to melt a substance = (mass of substance) x (its heat of fusion)Example: How much heat is needed to melt 35 g of ice cubes at 0.0 C?Q = mCp∆T Heat = 35 g x 334 J/g = 11,690 J = 1.169 x 104 J

Heat of vaporization: is the heat needed to vaporize a liquid at its boiling point.

Heat required to vaporize a liquid = (mass of substance) x (its heat of vaporization)Example: How many joules of heat are needed to convert 120 g water to steam at 100 C?Heat = 125 g x 2260 J/g = 2.825 x 105 J

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Note: MP and FP are the same temperature.

Heating and cooling curves:Phase Changes are reversible.They occur in pairs – one is endothermic, the other is exothermic.

Triple Point – The point at which all six phase changes can take placeCritical Point – The point above which only the gas phase can exist

Energy in chemical reactions:All chemical reactions are associated will energy change.

Activation energy: is the minimum energy required to initiate a chemical reaction – the energy needed to break the bonds in the reactants.

Activation complex is the temporary intermediate formed between the reactants and the products before the products are made.

Reactions are classified according to their energy:

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Solid Liquid

GasCritical Point

Triple Point

melting

freezing

boiling

condensation

sublimation

deposition

Exothermic reaction: when the energy of products is lower than the energy of the reactants. Heat is given of by the reaction.Example of exothermic reaction: CH4 + 2 O2 à CO2 + 2H2O + 213 kcal of heat

Endothermic reaction: when the energy of products is higher than the energy of reactants. Heat is absorbed by the reaction.Example of endothermic reaction: H2 + I2 + 12 kcal of heat à 2HI

Rate of reaction: is how fast the reactants are used up or the products are formed. The rate measures the speed of the reaction.

Factors that affect the rate of a reaction:1) The amount of reactant : as the amount of reactant increases the rate increases.2) Temperature : as temperature increases, the rate increases.3) Catalyst : is a substance that will increase the rate with out being part of the reaction (Note: in

Biology Catalysts are called Enzymes).

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Chemical Equilibrium:The reaction reaches equilibrium when its forward rate equals its reverse rate.

N2 (g) + 3H2 (g) <---> 2NH3 (g)Changes in equilibrium: we can introduce changes in equilibrium by:

1) Change in the concentrations (amounts) of reactants or products: According to La Chatelier’s principle; when a stress is applied to a reaction at equilibrium by changing the amount of reactants or products, then the rates of forward and reverse reactions will change to relieve that stress.

2) Change in Temperature: depends on whether the reaction is Exothermic or Endothermic;Exothermic reaction: increase in Temperature will shift the reaction to left.Endothermic reaction: increase in Temperature will shift the reaction to right.

Colloids, Suspensions and SolutionsColloids: are heterogeneous (non-uniform) mixtures that do not separate or settle out. Suspensions: are heterogeneous (non-uniform) mixtures that have larger particles than colloidal

particles and separate over time.Solutions: are homogeneous (uniform) mixtures of one substance (solute) dispersed in another solvent.

Solubility: Like dissolve like: the polarities of a solute and a solvent must be similar in order to form a solution.Example: -oil (nonpolar) does not dissolve in water (polar)

-Ethanol (polar) dissolve in water (polar).Water as a Solvent: water is polar molecule that forms with solutes what is called a hydrogen bond.Hydrogen bond: intermolecular force that occurs between molecules where partially positive hydrogen is attracted to the strongly electronegative atoms of O, N, or F in other molecules.

Formation of Solutions: Solutions are formed by a process called hydration (ions of solute are surrounded by water molecules)

Solubility and saturated SolutionsSolubility: is the amount in grams of solute dissolved in 100 g of solvent.Unsaturated solution: when a solution can dissolve more solute (a point under the line of a solubility

curve).Saturated solution: when a solution contains all the solute that it can dissolve (a point on the line of a

solubility curve).Supersaturated solution: when a solution has been prepared such that it has more solute dissolved in it

than a saturated solution (a point above the line of a solubility curve).

Effect of Temperature on Solubility:The solubility of most solids in water increases as temperature increases. While the solubility of gases

decreases as temperature increases.

Henry’s Law: the solubility of gas in a liquid increases as the pressure of that gas above the liquid increases.

Electrolytes:1) Strong electrolyte : a substance that when dissolves in water it separate completely into ions and

its solution conducts electricity very well. Like; NaCl, CaF2,..2) Weak electrolyte : a substance that when dissolves in water it separates slightly into ions and its

solution conducts electricity weakly. Like; HF, CuOH,..3) Nonelectrolyte : a covalent substance that when dissolves in water it does NOT separate into ions

and its solution does NOT conduct electricity. Like; sugar.

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Percent Concentration: is classified as;1) Mass% = [Mass of solute (g)/Mass of solution (g)] x 100

Example: What is the mass % of a solution prepared by dissolving 25 g NaCl in 115 g of water?Mass of solution = 25 + 115 = 140 g, Mass % = [25g NaCl/140.g solution] = 18%

2) Volume % = [Volume of solute (mL)/Volume of solution (mL)] x 100Example: What is the Volume % of a solution prepared by mixing 15 mL ethanol in 135 mL of water?Volume of solution = 15 + 135 = 150 mL, Volume % = [15 mL ethanol/150. mL solution] = 10.0%

3) Mass/Volume % = [Mass of solute (g)/Volume of solution (mL)] x 100Example: What is the Mass/Volume % of a solution prepared by dissolving 5.0 g KCl in 115 mL of solution?Mass/Volume % = [5.0 g KCl/115 mL solution] = 4.3%

Percent Concentrations as Conversion Factors:Example: How many grams of sucrose must be dissolved in 1.5 L of water to make 4.0 % (m/v)

solution?1.5 L (1000 mL/1L)(4.0g sucrose/100 mL solution) = 60. g of sucrose.

Molarity and DilutionMolarity (M): is the concentration measured as the number of moles of solute in 1.0 Liter solution. The

units of molarity are (moles/L) Molarity = moles of solute/Liters of solution. Molarity = Mass of solute/MW x Liters of solution. Example: What is the molarity of 45 g of NaCl in 0.85 L of solution? Molarity = [45 g NaCl/58.4 g/mole x 0.85 L] = 0.91

Molarity as a Conversion Factor:Example: How many grams of KCl are required to prepare 0.55 L of 1.5 M KCl solution?

Mass of KCl = (1.5 mole/L)(74.5 g/mole)(0.55 L) = 61 gExample: What is the volume of a 2.5 M solution of HCl that contains 3.2 mole of HCl? Volume = (3.2 mole/2.4 mole/L) = 1.3 L

Dilution: Is the process of preparing solutions from liquid solutes. (C1)(V1) = (C2)(V2) where C is the concentration and V is the volumeThe concentration can be expressed as % or as Molarity. Example: what is the new concentration (m/v%) when water is added to 45 mL of 15% (m/v) NaOH to make 650 mL of diluted NaOH solution? (C1)(V1) = (C2)(V2)(15%)(45ml) = (650ml)(C2)C2 = 15 x 45) / 650C2 = 1.03%Example: what volume of a 0.54 M HCl solution can be prepared by diluting 45 mL of a 1.2 M HCl solution? (M1)(V1) = (M2)(V2)(1.2M)(45ml) = (.54M)(V2)V2 = (1.2 M x 45 mL/0.54 M) = 100. mL

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Acids and BasesArrhenius definition of an acid: is the substance that produces hydrogen ion (H+) when it dissolves in

water. HCl + H2O à H+ + Cl-

Naming acids: are often named by writing hydro followed by the root of the name of the halogen, then -ic, and finally acid, HCl is hydrochloric acid, HF is hydrofluoric acid.HNO3 however is nitric acid, H2SO4 is sulfuric acid.

Arrhenius definition of a base: is the substance that dissociate to metal ion and hydroxide ion (OH-) when it dissolve in water. NaOH + H2O à Na+ + OH-

Naming Bases: are called metal hydroxide, NaOH is sodium hydroxide.Bronsted-Lowry acids and Bases:They defined acids and bases based on proton (H+) transfer.

An acid: is the substance that donates a proton (H+) in the reaction. A base: is the substance that accepts a proton (H+) in the reaction. HCl (acid) + H2O (base) à H3O+ (hydronium ion) + Cl-

NH3 (base) + H2O (acid) à NH4+(ammonium ion) + OH-

Conjugate acid-base pairs:a) Conjugate base of an acid: is formed when the acid loses H+Examples: acid Conjugate base

HCl Cl-

H2O OH-

NH4+ NH3

b) Conjugate acid of a base: is formed when the base accept a H+Examples: Base Conjugate acid

Cl- HClOH- H2ONH3 NH4

+

Strengths of Acids and Bases: Strong acids : They dissociate completely in water to form H3O+ (H+)

HCl + H2O à H3O+ + Cl-

Weak acids : They dissociate slightly in water to form H3O+ (H+)HCN + H2O ↔ H3O+ + CN-

Strong bases : They dissociate completely in water to form OH-

NaOH + H2O à Na+ + OH-

Weak bases : They dissociate slightly in water to form OH-

NH3 + H2O ↔ NH4+ + OH-

Ionization of water: water ionizes slightly to hydronium ion and hydroxide ion; H2O + H2O ↔ H3O+ + OH- OR: H2O ↔ H+ + OH-

The equilibrium constant for the ionization of water is called the ion-product of water “Kw”, where Kw = [H+][OH-]= 1.0 x 10-14

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Solutions are classified according to the concentration of H+ and OH-;1) Neutral solution : when [H+] = [OH-] = 1.0 x 10-7

2) Acidic solution : when [H+] > [OH-] 3) Basic solution : when [H+] < [OH-]

The pH scale: is a measure of the acidity of solutions. pH is a mathematical way to express the concentration of H+ or OH-

pH = -log[H+] and the range of pH is 0-14Solutions are classified according to their pH:1) Neutral solution: has pH=7 or [H+] = 1.0 x 10-7

2) Acidic solution: has pH<7 or [H+] > 1.0 x 10-7

3) Basic solution: has pH>7 or [H+] < 1.0 x 10-7

Example 1: Calculate the pH and determine whether the solution as acidic, basic, or neutral if [H+] = 1.0 x 10-5 MpH = - log[1.0 x 10-5] = 5.00, and it is acidic. Example 2: Calculate the pH of bleach that has [OH-]= 2.00 x 10-3 MpOH = -log(2.00 x 10-3) = 2.7pH + pOH = 14 pH = 14 – 2.7 = 11.3Example 3: Calculate the [H+], if the pH = 3[H+] = 1.0 x 10-pH [H+] = 1.0 x 10-3.

Reaction of Acids and Bases:1) Reaction of metals with acid:

Metal + Acid à Salt + WaterZn + HCl à ZnCl2 + H2O

2) Reaction of acids with carbonate and bicarbonate: HCl + NaHCO3 à CO2 + H2O + NaCl2HCl + Na2CO3 à CO2 + H2O + 2NaCl

3) Reaction of acid with hydroxide: is called “neutralization” Neutralization: is the reaction between acid and base to form salt and water. Acid + Base à Salt + WaterHCl + NaOH à NaCl + H2O

Buffer solution: is a solution that resists the change in the pH when small amounts of acid or base are added. Buffer solution is made of a combination of a weak acid and a salt containing its conjugate base, like blood buffer that is made of H2CO3(carbonic acid) and HCO3

- (bicarbonate) CO2 + H2O ↔ H2CO3 ↔ H3O+ + HCO3

-

Also buffers can be made of a combination of a base and a salt containing its conjugate acid.

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