Chemistry Unit 4 PPT

Chemistry
Toombs County High School

Energy is related to
molecular motion,
endothermic and exothermic
processes, and the states of
matter interrelated in
chemical reactions.

How does the quantity
of energy impact the
rate of reactions?

SC2.Students will relate how the
Law of Conservation of Matter is
used to determine chemical
composition in compounds and
chemical reactions.

c) Apply concepts of the mole
and Avogadro’s number to
conceptualize and calculate
molar volumes of gases.

d) Identify and solve
different types of
stoichiometry problems,
specifically relating mass to
moles and mass to mass.

e) Demonstrate the
conceptual principle of
limiting reactants.

SC2 Students will relate how the
chemical reactions.

 If a carpenter had two tabletops
and seven table legs, he could
only build one four-legged
table. The number of table legs
is the limiting factor in the
construction of four-legged
tables. Similarly, in chemistry,
the amount of product made in
a chemical reaction may be
limited by the amount of one or
more of the reactants.
12.3

Limiting and Excess Reagents
How is the amount of product in a
reaction affected by an insufficient
quantity of any of the reactants?
12.3

▪In a chemical reaction, an insufficient
quantity of any of the reactants will
limit the amount of product that
forms.
▪The limiting reagent is the reagent that
determines the amount of product that
can be formed by a reaction.
12.3

▪In the reaction of nitrogen and
hydrogen, hydrogen is the limiting
reagent. Nitrogen is the reagent that
is not completely used up in the
reaction. The reagent that is not used
up is called the excess reagent.
12.3

12.3
The Chemical Equation for the Preparation of
Ammonia

Problem Solving 12.25 Solve
Problem 25 with the help of an
interactive guided tutorial.

 Percent Yield
 What does the percent yield of a reaction
measure?
12.3

▪ The percent yield is a measure of the efficiency
of a reaction carried out in the laboratory.
▪ A batting average is actually a percent yield.
12.3

▪The theoretical yield is the maximum
amount of product that could be formed
from given amounts of reactants.
▪In contrast, the amount of product that
actually forms when the reaction is
carried out in the laboratory is called the
actual yield.
12.3

▪ The percent yield is the ratio of the
actual yield to the theoretical yield
expressed as a percent.
12.3

SC7.Students will characterize
the properties that describe
solutions and the nature of acids
and bases.

a) Explain the process of dissolving
in terms of solute/solvent
interactions: Observe factors that
affect the rate at which a solute
dissolves in a specific solvent,
Relate molality to colligative
properties.

 A sinkhole forms when the
roof of a cave weakens from
being dissolved by
groundwater and suddenly
collapses. One recorded
sinkhole swallowed a house,
several other buildings, five
cars, and a swimming pool!
You will learn how the
solution process occurs and
16.1

 Solution Formation
 What factors determine the rate at which a
substance dissolves?
16.1

▪ The compositions of the solvent and the solute
determine whether a substance will dissolve. The factors
that determine how fast a substance dissolves are
stirring (agitation)
temperature
the surface area of the dissolving particles
16.1

▪ A cube of sugar in cold tea dissolves slowly.
16.1

▪ Granulated sugar dissolves in cold water more quickly
than a sugar cube, especially with stirring.
16.1

▪ Granulated sugar dissolves very quickly in hot tea.
16.1

 Stirring and Solution Formation
▪ Stirring speeds up the dissolving process because fresh
solvent (the water in tea) is continually brought into
contact with the surface of the solute (sugar).
16.1

 Temperature and Solution Formation
▪ At higher temperatures, the kinetic energy of water
molecules is greater than at lower temperatures, so they
move faster. As a result, the solvent molecules collide
with the surface of the sugar crystals more frequently
and with more force.
16.1

 Particle Size and Solution Formation
▪ A spoonful of granulated sugar dissolves more quickly
than a sugar cube because the smaller particles in
granulated sugar expose a much greater surface area to
the colliding water molecules.
16.1

 Solubility
 How is solubility usually expressed?
16.1

▪ A saturated solution contains the maximum amount of
solute for a given quantity of solvent at a given
temperature and pressure.
▪ An unsaturated solution contains less solute than a
saturated solution at a given temperature and pressure.
16.1

▪ In a saturated solution, the
rate of dissolving equals the
rate of crystallization, so
the total amount of
dissolved solute remains
constant.
16.1

▪ The solubility of a substance is the amount of solute that
dissolves in a given quantity of a solvent at a specified
temperature and pressure to produce a saturated solution.
▪ Solubility is often expressed in grams of solute per 100 g of
solvent.
16.1

▪ Some liquids combine in all proportions, while others
don’t mix at all.
▪ Two liquids are miscible if they dissolve in each other in all
proportions.
▪ Two liquids are immiscible if they are insoluble in each other.
16.1

▪ Oil and water are immiscible.
16.1

▪ Vinegar and oil are immiscible.
16.1

 Factors Affecting Solubility
 What conditions determine the amount of solute
that will dissolve in a given solvent?
16.1

 Temperature affects the solubility of solid, liquid,
and gaseous solutes in a solvent; both
temperature and pressure affect the solubility of
gaseous solutes.
16.1

 Temperature
▪ The solubility of most solid substances increases as the
temperature of the solvent increases.
▪ The solubilities of most gases are greater in cold water than in
hot.
16.1

▪ The mineral deposits
around hot springs result
from the cooling of the
hot, saturated solution of
minerals emerging from
the spring.
16.1

▪ A supersaturated solution contains more solute than it
can theoretically hold at a given temperature.
▪ The crystallization of a supersaturated solution can be
initiated if a very small crystal, called a seed crystal, of
the solute is added.
16.1

▪ A supersaturated solution is clear before a seed crystal is
added.
16.1

▪ Crystals begin to form in the solution immediately after
the addition of a seed crystal.
16.1

▪ Excess solute crystallizes rapidly.
16.1

 Simulation 20
 Observe the effect of temperature on the
solubility of solids and gases in water.

 Pressure
▪ Changes in pressure have little effect on the solubility of
solids and liquids, but pressure strongly influences the
solubility of gases.
▪ Gas solubility increases as the partial pressure of the gas
above the solution increases.
16.1

▪ Henry’s law states that at a given temperature, the
solubility (S) of a gas in a liquid is directly proportional to
the pressure (P) of the gas above the liquid.
16.1

Problem Solving 16.2 Solve Problem 2
with the help of an interactive guided
tutorial.

 1. For a given substance, which of the following
will NOT influence how fast it dissolves?
▪ temperature
▪ amount of agitation
▪ molar mass
▪ size of the crystals

 2. The solubility of a substance is often expressed
as the number of grams of solute per
▪ 100 liters of solvent.
▪ 1 cm3
of solvent.
▪ 100 grams of solution.
▪ 100 grams of solvent.

 3. The solubility of a gas in a solvent is affected by
▪ both temperature and pressure.
▪ only pressure.
▪ only temperature.
▪ both pressure and agitation.

 Water must be tested
continually to ensure that the
concentrations of contaminants
do not exceed established
limits. These contaminants
include metals, pesticides,
bacteria, and even the by-
products of water treatment.
You will learn how solution
concentrations are calculated.
16.2

16.2
 Molarity
 How do you calculate the molarity of a solution?

▪ The concentration of a solution is a measure of the
amount of solute that is dissolved in a given quantity of
solvent.
▪ A dilute solution is one that contains a small amount of solute.
▪ A concentrated solution contains a large amount of solute.
16.2

▪ Molarity (M) is the number of moles of solute dissolved
in one liter of solution.
▪ To calculate the molarity of a solution, divide the
moles of solute by the volume of the solution.
16.2

▪ To make a 0.5-molar (0.5M) solution, first add 0.5 mol of
solute to a 1-L volumetric flask half filled with distilled
water.
16.2

▪ Swirl the flask carefully to dissolve the solute.
16.2

▪ Fill the flask with water exactly to the 1-L mark.
16.2

tutorial.
for Sample Problem 16.2

ProblemSolving 16.11 Solve Problem 11
tutorial.

16.2
 Making Dilutions
 What effect does dilution have on the total moles
of solute in a solution?

 Diluting a solution reduces the number of moles
of solute per unit volume, but the total number of
moles of solute in solution does not change.
16.2

▪ The total number of moles of solute remains unchanged
upon dilution, so you can write this equation.
▪ M1 and V1 are the molarity and volume of the initial solution,
and M2 and V2 are the molarity and volume of the diluted
solution.
16.2

▪ Making a Dilute Solution
16.2

▪ To prepare 100 ml of 0.40M MgSO4 from a stock solution
of 2.0M MgSO4, a student first measures 20 mL of the
stock solution with a 20-mL pipet.
16.2

▪ She then transfers the 20 mL to a 100-mL volumetric
flask.
16.2

▪ Finally she carefully adds water to the mark to make 100
mL of solution.
16.2

▪ Volume-Measuring Devices
16.2

tutorial.
`

 Percent Solutions
 What are two ways to express the percent
concentration of a solution?
16.2

▪ The concentration of a solution in percent can be expressed in
two ways: as the ratio of the volume of the solute to the volume
of the solution or as the ratio of the mass of the solute to the
mass of the solution.
16.2

 Concentration in Percent (Volume/Volume)
16.2

▪ Isopropyl alcohol (2-propanol) is sold as a 91% solution.
This solution consist of 91 mL of isopropyl alcohol mixed
with enough water to make 100 mL of solution.
16.2

Problem-Solving 16.15 Solve Problem 15
tutorial.

 Concentration in Percent (Mass/Mass)
16.2

 1. To make a 1.00M aqueous solution of NaCl,
58.4 g of NaCl are dissolved in
▪ 1.00 liter of water.
▪ enough water to make 1.00 liter of solution
▪ 1.00 kg of water.
▪ 100 mL of water.

 2. What mass of sodium iodide (NaI) is contained in
250 mL of a 0.500M solution?
▪ 150 g
▪ 75.0 g
▪ 18.7 g
▪ 0.50 g

 3. Diluting a solution does NOT change which of
the following?
▪ concentration
▪ volume
▪ milliliters of solvent
▪ moles of solute

 4. In a 2000 g solution of glucose that is labeled
5.0% (m/m), the mass of water is
▪ 2000 g.
▪ 100 g.
▪ 1995 g.
▪ 1900 g.

16.4
 Molality and Mole Fraction
 What are two ways of expressing the
concentration of a solution?

 The unit molality and mole fractions are two
additional ways in which chemists express the
concentration of a solution.
16.4

▪ The unit molality (m) is the number of moles of solute
dissolved in 1 kilogram (1000 g) of solvent. Molality is
also known as molal concentration.
16.4

▪ To make a 0.500m solution
of NaCl, use a balance to
measure 1.000 kg of water
and add 0.500 mol (29.3 g)
of NaCl.
16.4

▪ Ethlylene Glycol (EG) is added to water as antifreeze.
16.4

tutorial.

SC7.Students will
characterize the properties
that describe solutions and
the nature of acids and
bases.

b) Compare, contrast, and
evaluate the nature of acids
and bases: Strong vs. weak
acids/bases in terms of
percent dissociation
Bronsted-Lowery pH.

 Bracken Cave, near San
Antonio, Texas, is home to
twenty to forty million bats.
Visitors to the cave must
protect themselves from
the dangerous levels of
ammonia in the cave.
Ammonia is a byproduct of
the bats’ urine. You will
learn why ammonia is
19.1

 Properties of Acids and Bases
 What are the properties of acids and bases?
19.1

 Acids
▪ Acids taste sour, will change the color of an acid-base
indicator, and can be strong or weak electrolytes in
aqueous solution.
19.1

▪ Citrus fruits contain citric acid. Tea contains tannic acid.
19.1

 Bases
▪ Bases taste bitter, feel slippery, will change the color of
an acid-base indicator, and can be strong or weak
electrolytes in aqueous solution.
19.1

▪ Antacids use bases to neutralize excess stomach acid.
The base calcium hydroxide is a component of mortar.
19.1

 Arrhenius Acids and Bases
▪ How did Arrhenius define an acid and a base?
19.1

 Arrhenius said that acids are hydrogen-containing
compounds that ionize to yield hydrogen ions (H+
)
in aqueous solution. He also said that bases are
compounds that ionize to yield hydroxide ions
(OH–
) in aqueous solution.
19.1

 Arrhenius Acids
▪ Acids that contain one ionizable hydrogen, such as nitric acid
(HNO3), are called monoprotic acids.
▪ Acids that contain two ionizable hydrogens, such as sulfuric acid
(H2SO4), are called diprotic acids.
▪ Acids that contain three ionizable hydrogens, such as phosphoric
acid (H3PO4) are called triprotic acids.
19.1

 Arrhenius Bases
▪ Hydroxide ions are one of the products of the dissolution
of an alkali metal in water.
19.1

 Milk of magnesia is a base used as an antacid.
19.1

 Brønsted-Lowry Acids and Bases
 What distinguishes an acid from a base in the
Brønsted-Lowry theory?
19.1

 The Brønsted-Lowry theory defines an acid as a
hydrogen-ion donor, and a base as a hydrogen-ion
acceptor.
19.1

 Why Ammonia is a Base
19.1

 Conjugate Acids and Bases
▪ A conjugate acid is the particle formed when a base gains a
hydrogen ion.
▪ A conjugate base is the particle that remains when an acid has
donated a hydrogen ion.
19.1

▪ A conjugate acid-base pair consists of two substances related by
the loss or gain of a single hydrogen ion.
▪ A substance that can act as both an acid and a base is said to be
amphoteric.
19.1

▪ A water molecule that gains a hydrogen ion becomes a
positively charged hydronium ion (H3O+
).
19.1

 Lewis Acids and Bases
 How did Lewis define an acid and a base?
19.1

 Lewis proposed that an acid accepts a pair of
electrons during a reaction, while a base donates a
pair of electrons.
19.1

▪ A Lewis acid is a substance that can accept a pair of electrons to
form a covalent bond.
▪ A Lewis base is a substance that can donate a pair of electrons to
form a covalent bond.
19.1

 Animation 25
 Compare the three important definitions of acids
and bases.

 Problem Solving 19.1
 Solve Problem 1 with the help of an

 1. Which of the following is NOT a characteristic
of acids?
▪ taste sour
▪ are electrolytes
▪ feel slippery
▪ affect the color of indicators

 2. Which compound is most likely to act as an
Arrhenius acid?
▪ H2O
▪ NH3.
▪ NaOH.
▪ H2SO4.

 3. A Lewis acid is any substance that can accept
▪ a hydronium ion.
▪ a proton.
▪ hydrogen.
▪ a pair of electrons.

 To test a diagnosis of
diabetic coma, a doctor
orders several tests,
including the acidity of
the patient’s blood.
Results from this test will
be expressed in units of
pH. You will learn how the
pH scale is used to
indicate the acidity of a
19.2

 Hydrogen Ions from Water
▪ The reaction in which water molecules produce ions is
called the self-ionization of water.
19.2

▪ In the self-ionization of water, a proton (hydrogen ion)
transfers from one water molecule to another water
molecule.
19.2

 Ion Product Constant for Water
 How are [H+
] and [OH-
] related in an aqueous
solution?
19.2

 For aqueous solutions, the product of the
hydrogen-ion concentration and the
hydroxide-ion concentration equals 1.0 ×
10-14
.
 Any aqueous solution in which [H+
] and
[OH-
] are equal is described as a neutral
solution.
19.2

▪ The product of the concentrations of the hydrogen ions
and hydroxide ions in water is called the ion-product
constant for water (Kw).
19.2

▪ An acidic solution is one in which [H+
] is greater than
[OH-
].
19.2

▪ Unrefined hydrochloric acid, commonly called muriatic
acid, is used to clean stone buildings and swimming
pools.
19.2

▪ A basic solution is one in which [H+
] is less than [OH−
].
Basic solutions are also known as alkaline solutions.
19.2

▪ Sodium hydroxide, or lye, is commonly used as a drain
cleaner.
19.2

 The pH Concept
 How is the hydrogen-ion concentration used to
classify a solution as neutral, acidic, or basic?
19.2

▪ The pH of a solution is the negative logarithm of the
hydrogen-ion concentration.
19.2

▪ A solution in which [H+
] is greater than 1 × 10–7
M has a pH
less than 7.0 and is acidic. The pH of pure water or a
neutral aqueous solution is 7.0. A solution with a pH
greater than 7 is basic and has a [H+
] of less than 1 × 10–7
M.
19.2

 pH and Significant Figures
19.2

 Solve Problem 14 with the help of
an interactive guided tutorial.

 Measuring pH
 What is the most important characteristic of an
acid-base indicator?
19.2

 An indicator is a valuable tool for measuring pH
because its acid form and base form have
different colors in solution.
19.2

▪ Phenolphthalein changes from colorless to pink at pH 7–
9.
19.2

 1. If the [OH-
] in a solution is 7.65 × 10-3
M, what is
the [H+
] of this solution?
▪ 7.65 × 10-17
M
▪ 1.31 × 10-12
M
▪ 2.12M
▪ 11.88M

 2. The [OH-
] for four solutions is given below.
Which one of the solution is basic?
▪ 1.0 x 10-6
M
▪ 1.0 x 10-8
M
▪ 1.0 x 10-7
M
▪ 1.0 x 10-14
M

 3. What is the pH of a solution with a hydrogen-
ion concentration of 8.5 x 10-2
M?
▪ 12.93
▪ 8.50
▪ 5.50
▪ 1.07

 Lemons and grapefruits
have a sour taste because
they contain citric acid.
Sulfuric acid is a widely
used industrial chemical
that can quickly cause
severe burns if it comes
into contact with skin.
You will learn why some
acids are weak and some
19.3

 Strong and Weak Acids and Bases
 How does the value of an acid dissociation
constant relate to the strength of an acid?
19.3

▪ An acid dissociation constant (Ka) is the ratio of the
concentration of the dissociated (or ionized) form of an
acid to the concentration of the undissociated
(nonionized) form.
19.3

 Weak acids have small Ka values. The stronger an
acid is, the larger is its Ka value.
19.3

▪ Strong acids are completely ionized in aqueous
solution.
▪ Weak acids ionize only slightly in aqueous solution.
19.3

▪ In general, the base dissociation constant (Kb) is the
ratio of the concentration of the conjugate acid times
the concentration of the hydroxide ion to the
concentration of the base.
19.3

▪ Strong bases dissociate completely into metal ions and
hydroxide ions in aqueous solution.
▪ Weak bases react with water to form the hydroxide ion
and the conjugate acid of the base.
19.3

 Calculating Dissociation Constants
 How can you calculate an acid dissociation
constant (Ka) of a weak acid?
19.3

 To find the Ka of a weak acid or the Kb of a weak
base, substitute the measured concentrations of
all the substances present at equilibrium into the
expression for Ka or Kb.
19.3

 Acid Dissociation Constant
▪ The dissociation constant, Ka, of ethanoic acid is
calculated from the equilibrium concentrations of all of
the molecules and ions in the solution.
19.3

 Base Dissociation Constant
▪ The dissociation constant, Kb, of ammonia is calculated
from the equilibrium concentrations of all of the
molecules and ions in the solution.
19.3

 Concentration and Strength
19.3

 Solve Problem 23 with the
help of an interactive guided
tutorial.

 1. H2S is considered to be a weak acid because it
▪ is insoluble in water.
▪ ionizes only slightly.
▪ is completely ionized.
▪ is dilute.

 2. Calcium hydroxide, Ca(OH)2, is a strong base
because it
▪ has a large Kb.
▪ has a small Kb.
▪ forms concentrated solutions.
▪ is highly soluble in water.

 3. If the [H+
] of a 0.205M solution of phenol (C6H5OH)
at 25ºC is 2.340 × 10-6
, what is the Ka for phenol?
Phenol is monoprotic.
▪ Ka = 2.67 x 10-11
▪ Ka = 1.14 x 10-5
▪ Ka = 5.48 x 10-12
▪ Ka = 1.53 x 10-3

 4. The Ka of three acids is given below.
 (1) 5.1 × 10–3
 (2) 4.8 × 10–11
 (3) 6.3 × 10–5
 Put the acids in order from the strongest acid to the weakest
acid.
▪ 1, 3, 2
▪ 2, 3, 1
▪ 3, 1, 2
▪ 2, 1, 3

 5. The Kb of four bases is given below.
 (1) 7.41 x 10-5
 (2) 1.78 x 10-5
 (3) 4.27 x 10-4
 (4) 4.79 x 10-4
 Put the bases in order from the strongest base to the weakest base.
▪ 2, 3, 4, 1
▪ 2, 1, 3, 4
▪ 4, 3, 1, 2
▪ 1, 4, 3, 2

 Excess hydrochloric acid in the stomach can
cause heartburn and a feeling of nausea.
Antacids neutralize the stomach acid and
relieve the pain of acid indigestion. You will
learn what a neutralization reaction is.
19.4

 Acid-Base Reactions
 What are the products of the reaction of an acid
with a base?
19.4

▪ In general, the reaction of an acid with a base produces
water and one of a class of compounds called salts.
19.4

▪ Reactions in which an acid and a base react in an
aqueous solution to produce a salt and water are
generally called neutralization reactions.
19.4

 Titration
 What is the endpoint of a titration?
19.4

▪ The process of adding a known amount of solution of
known concentration to determine the concentration of
another solution is called titration.
▪ The point of neutralization is the end point of the titration.
19.4

▪ When an acid and base are mixed, the equivalence
point is when the number of moles of hydrogen ions
equals the number of moles of hydroxide ions.
19.4

 Solve Problem 30 with the help
of an interactive guided tutorial.

▪ The solution of known concentration is called the
standard solution.
▪ Indicators are often used to determine when enough of the
standard solution has been added to neutralize the acid or base.
▪ The point at which the indicator changes color is the end point of
the titration.
19.4

Acid solution with
indicator
Added base is
measured with a buret.
Color change
shows
neutralization.
19.4

 Simulation 26
 Simulate the titration of several acids and bases
and observe patterns in the pH at equivalence.

 Solve Problem 33 with the help
of an interactive guided tutorial.

 1. When a neutralization takes place, one of the
products is always
▪ carbon dioxide.
▪ a salt.
▪ sodium chloride.
▪ a precipitate.

 2. In a titration, 45.0 mL of KOH is neutralized by
75.0 mL of 0.30M HBr. What is the concentration
of the KOH solution?
▪ 0.18M
▪ 0.60M
▪ 0.25M
▪ 0.50M

 3. How many moles of HCl are required to
neutralize an aqueous solution of 2.0 mol
Ca(OH)2?
▪ 0.5 mol
▪ 1.0 mol
▪ 2.0 mol
▪ 4.0 mol

 4. In which of the following neutralization
titrations of 1-molar solutions of H2SO4 and NaOH
will the equivalence point be reached at the very
end of the additions?
 H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(aq)
▪ 200 mL of H2SO4 is slowly added to 100 mL of NaOH

 The chemical processes inside a living cell are
very sensitive to pH. Human blood is
normally maintained at a pH very close to 7.4.
You will learn about chemical processes that
ensure that the pH of blood is kept near 7.4.
19.5

 Salt Hydrolysis
 When is the solution of a salt acidic or basic?
19.5

 In general, salts that produce acidic solutions
contain positive ions that release protons to
water. Salts that produce basic solutions contain
negative ions that attract protons from water.
19.5

▪ In salt hydrolysis, the cations or anions of a dissociated
salt remove hydrogen ions from or donate hydrogen
ions to water.
19.5

▪ To determine whether a salt solution is acidic or basic,
remember the following rules:
19.5

▪ Vapors of the strong acid HCl(aq) and the weak base
NH3(aq) combine to form the acidic white salt
ammonium chloride (NH4Cl).
19.5

NH4Cl
pH 5.3
NaCl
pH 7
CH3COONa
pH 5.3
▪ Universal indicator solution has been added to each of
these 0.10M aqueous salt solutions.
19.5

 Buffers
 What are the components of a buffer?
19.5

▪ A buffer is a solution of a weak acid and one of its salts, or
a solution of a weak base and one of its salts.
The pH of a buffer remains relatively constant
when small amounts of acid or base are
added.
The buffer capacity is the amount of acid or
base that can be added to a buffer solution
before a significant change in pH occurs.
19.5

▪ Buffer of Ethanoic Acid and Sodium Ethanoate
▪ Adding H+
produces additional ethanoic acid.
▪ Adding OH-
produces additional ethanoate ions.
▪ The pH changes very little.
19.5

 Animation 26
 Discover the chemistry behind buffer action.

 Solve Problem 39 with the help of
an interactive guided tutorial.

 1. Which of the following reactions would most
likely yield a basic salt solution?
▪ strong acid + weak base
▪ weak acid + weak base
▪ strong acid + strong base
▪ weak acid + strong base

 2. Choose the correct words for the spaces. A
buffer can be a solution of a _________ and its
_________.
▪ weak acid, salt
▪ strong acid, salt
▪ weak acid, conjugate base
▪ weak base, conjugate acid

 3. Which of the following equations represents the
reaction when a high pH substance is added to a
dihydrogen phosphate ion-hydrogen phosphate
ion buffer system?
▪ H2PO4¯+ OH¯→ HPO4
2
¯ + H2O
▪ HPO4
2
¯+ OH¯→ PO4
3
¯ + H2O
▪ H2PO4¯+ H+
→ H3PO4
▪ HPO4
2
¯+ H+
→ H2PO4¯

 Concept Map 19
 Create your Concept Map using the
computer.

SC6. Students will
understand the effects of
motion on atoms and
molecules in chemical and
physical processes.

a) Compare and contrast
atomic/molecular motion in
solids, liquids, gases, and
plasmas.

b) Collect data and
calculate the amount of
heat given off or taken in by
chemical or physical
processes.

c) Analyze (both
conceptually and
quantitatively) the flow of
energy during change of
state (phase).

SC2.Students will relate how the
chemical reactions.

f) Explain the role of
equilibrium in chemical
reactions.

SC6.Students will understand
the effects of motion on atoms
and molecules in chemical
and physical processes.

b) Collect data and calculate
the amount of heat given off
or taken in by chemical or
physical processes.

SC5.Students will understand that
the rate at which a chemical
reaction occurs can be affected by
changing concentration,
temperature, or pressure and the
addition of a catalyst.

a) Demonstrate the effects
of changing concentration,
temperature, and pressure
on chemical reactions.

b) Investigate the effects of a
catalyst on chemical
reactions and apply it to
everyday examples.

c) Explain the role of
activation energy and
degree of randomness in
chemical reactions.

Chemistry Unit 4 PPT

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