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Chemistry Minimum Learning Outcomes – The
student should 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 milliliter
3. 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) 1000 = 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 factors
Quantity (given unit) x conversion factor (desired unit/given unit) = desired unit
Example 1: How many moles are in 45 grams of sodium?
Plan: g  mole
45 g Na x 1 mole Na
= 1.96 mole Na
22.99 g Na
Example 2: A physician has ordered 325 mg of atropine, intramuscularly. If atropine were
available as 0.50 g/mL of solution, how many mL would you need to give?
Plan: mg  g  mL
325 mg x
1g
x
1 mL
= 0.65 mL
1000 mg
0.50 g
Density: is the mass per unit volume.
Density (D) = Mass(M)/volume(V), or [D=M/V]
*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/cm3
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 g
Example3: 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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Temperature
Temperature: 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 C
2. Kelvin: K
The relationship between the three scales of temperature:
K= C + 273 or C = K – 273
Examples: Convert 33.0 C to K.
answer: K = C + 273 = 33.0 + 273 = 306 K
Example: Convert 450. K to C.
answer: C= K – 273 = 450. – 273 = 177 C
Elements and symbols
Elements: 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 1A
2) Alkaline earth metals: are the elements in group 2A
3) Halogens: are the elements in group 7A
4) 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, amu
proton
in nucleus
+1
1
neutron
in nucleus
0
1
electron
in orbit around nucleus
-1
1/2000
Atomic number: the number of protons in the nucleus.
In the neutral atom the number of electrons equals the number of protons.
Mass number: 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 31H.
Example: identify the number of protons and neutrons in the 2512Mg 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 table
Example: 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 orbital, where each orbital can hold a maximum of two
electrons. There are four different kinds of orbitals, which are classified according to their
shapes to: 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?
Answer: 1s2 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 Bonds
Valance 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 helium and hydrogen they hold 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 to:
Positive ions (Cations): are atoms that have lost electron(s). Like; Li loses 1e- and become
Li+ ion, Ca loses 2e- and become Ca+2 ion, Al loses 3e- and become Al+3 ion,…
Negative ions (Anions): are atoms that have gained electron(s). Like; F gains 1e- and
become F- ion, O gains 2 e- and become O-2 ion, N gains 3 e- and 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 Compounds
Ionic 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 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 charge
Examples: 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
First write the ions:
Na+
HCO3Balance charges (crisscross):
NaHCO3
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aluminum sulfate
Al+3
SO4-2
AL2(SO4)3
magnesium phosphate
Mg+2
PO4-3
Mg3(PO4)2
Covalent Compounds
Covalent 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 : H
Double bond: when the two atoms share 2 pairs of electrons, O::O, or C::C
Triple 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
SO3
CCl4
silicon dioxide
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 molecules
To 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. Like; CH4, CO2
2) Polar molecule: When the molecule has only one polar bond or when polar bonds do not
cancel each other. Like; HCl, H2O.
Chemical reactions and quantities
Physical change: is when a physical property such as the appearance of a substance is altered,
but not its composition.
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  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  CO2 + 2 H2O
The 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  2NO2
2) Decomposition reaction: a reactant splits into tow or more simple products, example:
CaCO3  CaO + CO2
3) Single replacement reaction: Na + HCl  NaCl + H2
4) Double replacement reaction: BaCl2 + Na2SO4  BaSO4 + 2NaCl
5) Combustion reaction: fuel and oxygen react to produce carbon dioxide and water.
CH4 + 2O2  CO2 + H2O
Oxidation-Reduction Reactions (LEO GER)
Oxidation: is loss of electrons. Na  Na+ + eReduction: is gain of electrons. O + 2e-  O2Example: Ca + S  CaS in this reaction “Ca” is oxidized and “S” is reduced.
The Mole
The mole is a unit used to count the number of atoms, molecules, or formulas. For example: one
mole of carbon atoms contains 6.02 x 1023 atoms of Carbon.
One mole of hydrogen atoms in 1 mole hydrogen molecules (H2) is 2 x 6.02 x 1023
1 Mole = mass of element in grams / molecular weight or atomic weight.
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 equations
From 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 Matter
Energy 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 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: heat = (specific heat)(mass)(temperature change, ∆T)
Example:
Calorimeter: is a device used to measure the energy transfer between objects.
Example:
States of Matter: Solid, Liquid, Gas
Changes 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
Note: MP and FP are
the same temperature.
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?
Heat = 35 g x 334 J/g = 11,690 J = 1.169 x 104 J
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 boil.
Condensation: is the exothermic change of vapor to a liquid.
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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Heating and cooling curves:
Phase Changes are reversible.
They occur in pairs – one is endothermic, the other is exothermic.
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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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  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  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:
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.
Chemical Equilibrium:
The reaction reaches equilibrium when its forward rate equals its reverse rate.
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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.
Solutions
Solution: is a uniform mixture of one substance (solute) dispersed in another solvent.
Types of Solutes and Solvents:
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: 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 Solutions
Solubility: is the amount in grams of solute dissolved in 100 g of solvent.
Unsaturated solution: when a solution can dissolve more solute.
Saturated solution: when a solution contains all the solute that can dissolve.
Supersaturated solution: when a solution has been prepared such that it has more solute dissolved
in it than a saturated solution.
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 to ions completely
and its solution conduct electricity. Like; NaCl, CaF2,..
2) Weak electrolyte: a substance that when dissolves in water it separate to ions slightly and
its solution conduct electricity. Like; HF, CuOH,..
3) Nonelectrolyte: a covalent substance that when dissolves in water it does NOT separate
to ions and its solution does NOT conduct electricity. Like; sugar.
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Equivalents (Eq): is the amount of the ion that equals to 1.0 mole of positive or negative charge.
Some times is measured in milli-equivalent (mEq).
Examples: Indicate the number of equivalents in the following:
a) 1mole of Na+
b) 2 moles OHc) 3 moles CO3-2
a) 1mole Na+ (1 Eq Na+/1 mole Na) = 1 Eq Na+
b) 2 mole OH- (1 Eq OH-/1 mole OH-) = 2 Eq OHc) 3 mole CO3-2 (2 Eq CO3-2/1 mole CO3-2) = 6 Eq CO3-2
Percent Concentration: is classified to;
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 100
Example: 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 100
Example: 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 Dilution
Molarity (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 g
Example: 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 volume
The 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) / 650
C2 = 1.03%
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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
Colloids and Suspensions
Colloids: are homogeneous 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.
Acids and Bases
Arrhenius definition of an acid: is the substance that produces hydrogen ion (H+) when it
dissolves in water. HCl + H2O  H+ + ClNaming acids:
Arrhenius definition of a base: is the substance that dissociate to metal ion and hydroxide ion
(HO-) when it dissolve in water. NaOH + H2O  Na+ + OHNaming 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)  H3O+ (hydronium ion) + ClNH3 (base) + H2O (acid)  NH4+(ammonium ion) + OHConjugate acid-base pairs:
a) Conjugate base of an acid: is formed when the acid loses H+
Examples:
acid
Conjugate base
HCl
ClH2O
OH+
NH4
NH3
b) Conjugate acid of a base: is formed when the base accept a H+
Examples:
Base
Conjugate acid
Cl
HCl
OHH2O
NH3
NH4+
Strengths of Acids and Bases:
 Strong acids: They dissociate completely in water to form H3O+
HCl + H2O  H3O+ + Cl Weak acids: They dissociate slightly in water to form H3O+
HCN + H2O ↔ H3O+ + CN Strong bases: They dissociate completely in water to form OHNaOH + H2O  Na+ + OH Weak bases: They dissociate slightly in water to form OHNH3 + H2O ↔ NH4+ + OH-
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Ionization of water: water ionizes slightly to hydronium ion and hydroxide ion;
H2O + H2O ↔ H3O+ + OHThe equilibrium constant for the ionization of water is called the ion-product of water
“Kw”, where Kw = [H3O+][OH-]= 1.0 x 10-14
Solutions are classified according to the concentration of H3O+ and OH- to;
1)
Neutral solution: when [H3O+] = [OH-] = 1.0 x 10-7
2)
Acidic solution: when [H3O+] > [OH-]
3)
Basic solution: when [H3O+] < [OH-]
Example: What is the concentration of H3O+ in a Bleach solution if [OH-] = 1.0 x 10-2. Is the
solution acidic, basic, or neutral?
[H3O+][OH-] = 1.0 x 10-14
[H3O+] = (1.0 x 10-14)/(1.0 x 10-12) = 1.0 x 10-12 , solution is basic.
Example: What is the concentration of OH- in a vinegar solution if [H3O+] = 2.0 x 10-3. Is
the solution acidic, basic, or neutral?
[H3O+][OH-] = 1.0 x 10-14
[OH-] = (1.0 x 10-14)/(2.0 x 10-3) = 5.0 x 10-12 , solution is acidic.
The pH scale: is a measure of the acidity of solutions.
pH is a mathematical way to express the concentration of H3O+ or OHpH = -log[H3O+] and the range of pH is 0-14
Solutions are classified according to their pH to:
1)
Neutral solution: has pH=7 or [H3O+] = 7.0 x 10-7
2)
Acidic solution: has pH<7 or [H3O+] > 7.0 x 10-7
3)
Basic solution: has pH>7 or [H3O+] < 7.0 x 10-7
Example 1: Calculate the pH and determine whether the solution as acidic, basic, or neutral,
if [H3O+] = 1.0 x 10-5 M
pH = - 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 M
[H3O+] = (1.0 x 10-14)/(2.00 x 10-3)= 1.00 x 10-11
pH = -log(1.00 x 10-11) = 11.0
Example 3: Calculate the [H3O+], if the pH = 3
[H3O+] = 1.0 x 10-3, because the pH is the negative power of 10.
Reaction of Acids and Bases:
1) Reaction of metals with acid:
Metal + Acid  Salt + Water
Zn
+ HCl  ZnCl2 + H2O
2) Reaction of acids with carbonate and bicarbonate:
HCl + NaHCO3  CO2 + H2O + NaCl
2HCl + Na2CO3  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  Salt + Water
HCl + NaOH  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+ + HCO3Also buffers can be made of a combination of a base and a salt containing its conjugate acid.
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