Transcript Chapter 14 Review- Chemical Periodicity

Chapter 12 ReviewChemical Periodicity
Chapter 12 Review - definitions
electronegativity
 periods
 atomic radius
 ionization energy
 periodic table

Chapter 12 Review - definitions
alkali metals
 halogens
 noble gases
 alkaline earth metals
 groups

Chapter 12 Review
The modern periodic table is arranged in
order of increasing: ( atomic number )
 The elements in Groups 1A through 7A
and Group 0 make up the:

( representative elements )

What are characteristics of the noble
gases?
( outer s and p sublevels filled,
belong to Group 0, called inert gas )
Chapter 12 Review
What is the number of electrons in the
outermost occupied energy level of an
element in Group 5A? ( 5 )
 Which have the same number of
electrons in their outermost energy
levels? a) K, Ca, Rb, Sr b) N, P, As, Sb

(b)
Chapter 12 Review
An element that contains an electron in a
d sublevel is: a) Mg b) Fe ( b )
 The elements that contain electrons in
the f sublevels are referred to as:

( inner transition metals )

The outermost energy level configuration
of the element chlorine is:
( 3s23p5 )
Chapter 12 Review
The element with 8 electrons in its 3d
sublevel is: a) Ar b) Ni ( b )
 As you move down a group of the
periodic table, atomic size generally:

( increases )

The largest atom from among the
following is: a) Fr b) Rb ( a )
Chapter 12 Review
The smallest atom from the following is:
( Cl )
a) Cl b) Si
 As the number of electrons added to the
same principal energy level increases,
atomic size generally:

( decreases )

Removing one electron from a gaseous
atom forms a:
( 1+ ion )
Chapter 12 Review
Among the elements listed, which would
show the largest increase between the
2nd and 3rd ionization energies? a) Ca
b) Zn ( a )
 Among the following, which element has
the lowest IE? a) Cs b) I ( a )
 Among the following, which has the
highest 2nd IE? a) Na b) Cl ( a )

Chapter 12 Review
Which of the following are always larger
than the neutral atoms from which they
are formed?
a) cations b) anions ( b )
 The smallest particle from among the
following is: a) Li b) Li1+ ( b )
 Which is the least electronegative?
a)
S b) Cs ( b )

Fill in requested information from given:
Element
2s2
3s23p3
Period No. Group No. Group Name
(2)
(3)
3s23p6 ( 3 )
4s1
(4)
Symbol
( 2A )
( alkaline
earth metal)
( 5A )
(representative
element)
(P)
(0)
( noble gas )
( Ar )
( 1A )
( alkali metal )
(K)
( Sc )
( Br )
4s23d1
(4)
( 3B )
( transition
metal )
4s24p5
(4)
( 7A )
( halogen )
( Be )
Chapter 12 Review

Arrange the following: Li, C, K, F,
Cs ( Cs, K, Li, C, F )
 decreasing atomic size:
( Cs, K, Li, C, F )
 increasing
ioniztion energy:
( F, C, Li, K, Cs )
 decreasing
electronegativity:
Chapter 12 Review
Atom
Larger
Li, K
(K)
( Li )
(K)
C, F
(C)
(F)
(C)
Mg, Ca ( Ca )
O, S
(S)
Greater Ionization Energy Lower Electronegativity
( Mg )
( Ca )
(O)
(S)
Chapter 12 Review

Explain why atoms with high
ionization energies typically also
have high electronegativities.
( A high ionization energy indicates that
an atom has a tight hold on it’s electrons. A
high electronegativity indicates an ability
to attract additional electrons. They both
refer to a desire to have electrons. )
Chapter 17 Review
“Reaction Rates and Equilibrium”
Chapter 17 Review
 Energy
that is available to do work
Free energy
is called ____.
Reaction
Rate is defined as the number of

____
atoms, ions, or molecules that react
in a given time to form products.
 What is the name of the minimum
energy that colliding particles must
have in order to react? Activation Energy
Chapter 17Review
 A substance that interferes with a
inhibitor
catalyst is a(n) ____.
 What is the arrangement of atoms
Activated complex
at the peak of an energy barrier?
 At equilibrium, what is the rate of
production of reactants compared
with the rate of production of
products? The same
Chapter 17 Review
What is the equilibrium constant
expression for the following reaction:
[Carbon Dioxide]
C(s) + O2(g) ↔ CO2(g)
[Carbon] [Oxygen]
 Which of the following is true about the
a the reaction is
combustion of carbon: a)
spontaneous, or b) entropy decreases?
 The rate of a chemical reaction normally
increases
____ as reactant concentration
increases.

Chapter 17 Review
 Why
does a higher concentration
make a reaction faster? More collisions /time
 The amount of disorder in a system
entropy
is measured by its ____.
disorder
 Entropy measures ____
 For a complex reaction, the reaction
Elementary steps
progress curve has several ____.
Chapter 17 Review
What happens to a reaction at
equilibrium when more reactant is
added to the system? The Reaction shifts to the right
 Which reaction results in the greatest
increase in entropy:
b
a) A → B, or b) A → 2B
 If a reaction has an equilibrium constant
(Keq) just greater than 1, how do we
Reaction slightly favors
interpret that information? The
Products at equilibrium

A
Chapter 17 Review
Lowering the activation energy barrier
catalyst works by ____.
 If sulfur dioxide and oxygen can be
made into sulfur trioxide, what is the
reverse reaction? SO  S + O
 Which variable is NOT required to
calculate the Gibbs free-energy
change for a chemical reaction:
a) change in enthalpy, or
b) temperature in oC? b
3
2
Chapter 17 Review
 What
is the effect of adding more
water to the following equilibrium
reaction: CO2 + H2O ↔ H2CO3 Shift right
 In an endothermic reaction at
equilibrium, what is the effect of
raising the temperature? Shift Right
 The energy that is available to do
Free energy
work in a reaction is called ____.
Chapter 17 Review
 What
is the numerical value (+ or -)
of Gibbs free-energy change for a
negative
spontaneous reaction?
 Which of the following systems has
the highest entropy:
a) 10 mL of water at 50 oC, or
b) 10 mL of water at 100 oC?
b
Chapter 17 Review
What happens to a catalyst in a
reaction? The catalyst remains nchanged
 Write the rate law for the following
reaction: A + 2B → C + D Rate = k[A][B]
 An elementary reaction converts
One step
reactants to products in ____.
The energy needed to make reactants
 Activation energy is ____.
Become products
 Is the melting of ice at a temperature
above 0 oC:
a) endothermic, or b) exothermic? a

2
Chapter 17 Review
 Consider
the reaction:
N2(g) + 3H2(g) ↔ 2NH3(g) Shift right
What is the effect of decreasing the
volume on the contained gases?
release
 Spontaneous reactions always ____
free energy.
 Why does a higher temperature
cause a reaction to go faster?
Increases the likelyhood of a successful collision
Chapter 17 Review
 What
2 factors determine whether
or not a reaction is spontaneous?
Increasing entropy and decreasing enthalpy
 What
physical state of nitrogen has
the highest entropy? gas
 What is another name for the
catalysts in your body? enzyme
Chapter 17 Review
What is the order of the following
reaction: A + 2B → C + D 3 order
 Why does a catalyst cause a reaction to
Lower activation energy barrier
proceed faster?
 In an equilibrium reaction with a Keq of
1 x 108, the ____ are favored. products
 Which of the following explains why
melting of ice is spontaneous at room
temperature and pressure:
a) it is accompanied by an increase in
entropy, or b) it is accompanied by an
a
increase in energy?

rd
Chapter 17 Review
Keq of a reaction is 4 x 10-7. At
reactants are favored.
equilibrium, the ____
 Another name for the activated
Transition state
complex is ____.
 Which change would shift the
following reaction to the right:
4HCl(g) + O2(g) ↔ 2Cl2(g) + 2H2O(g)
a) decrease of pressure, or
b) increase of pressure? b
 The
 In
Chapter 17Review
which of these systems is the
entropy decreasing: b
a) salt dissolving in water, or
b) a liquid cooling?
 Given: 2NClO(g) ↔ 2NO(g) + Cl2(g)
An analysis of the equilibrium
mixture in a 1 L flask is: NClO = 1.6
mol; NO = 6.4 mol; Cl2 = 0.49 mol.
Calculate the value of Keq. 7.84
Chapter 17 Review
 In
a two-step reaction mechanism,
how many elementary reactions
occur? 2
 The Ksp of calcium hydroxide is
6.5 x 10-6. If the concentration is
0.0010 M Ca(OH)2, what is the
final concentration of the calcium
ion?
2.4x10
-3
Chapter 17 Review
A
mixture of hydrogen and iodine
are in equilibrium with hydrogen
iodide, as shown in the equation:
H2 + I2 ↔ 2HI
Calculate the concentration of HI
when the Keq is 1 x 105, the
equilibrium concentration of H2 is
0.04 M, and the equilibrium 6 x 10
concentration of I2 is 0.009 M.
-5
Chapter 17 Review
Chapter 18 Review
Acids and Bases
Chapter 18 Review - definitions
acidic solution
 conjugate acid-base pair
 amphoteric
 alkaline solution
 Kw

Chapter 18 Review - definitions
hydroxide ion
 neutral solution
 hydronium ion
 Ka
 triprotic acid

Chapter 18 Review

A solution in which the hydroxide ion
concentration is 1 x 10-5 M is:
( basic )

In a neutral solution, the [ H1+ ] is:
( equal to [ OH1- ] )
Chapter 18 Review
What are the products of the self
ionization of water? ( OH- and H+ )
 Which is most basic: a) [ H+ ] = 1 x
10-11, or b) [ OH- ] = 1 x 10-4? ( a )
 The formula of the hydrogen ion is
often written as: ( H+ )

Chapter 18 Review
What is the pH of a solution in which
the [ H+ ] = 1 x 10-12?
( 12.0 )
 What is the pH of a 0.01 M
hydrochloric acid solution? ( 2.0 )
 If Ka for H2CO3 is 4.3 x 10-7, it means
that H2CO3 is:

( a poor hydrogen-ion donor )
Chapter 18 Review

Which of the following pairs consist
of a weak acid and a strong base?
a) ethanoic acid, sodium hydroxide
or b) sulfuric acid, sodium hydroxide
(a)
Chapter 18 Review

In the reaction NH41+ + H2O  NH3
+ H3O1+, water is acting as a(n):
( Bronsted-Lowry base )

A solution with a pH of 5.0:
( has a hydroxide ion concentration of
1 x 10-9 M )
Chapter 18 Review
With solutions of strong acids and
strong bases, the word strong refers
( the degree of ionization )
to:
 Which of these is an Arrhenius base?
a) KOH, or b) NH3?

(a)
Chapter 18 Review

Ethanoic acid ionizes in water as:
1%
CH3COOH + H2O
-->
<--
CH3COO1- + H3O1+
99%
The ethanoate ion (CH3COO1-) is
therefore: ( a good hydrogen-ion acceptor )
Chapter 18 Review
Which acid is monoprotic?
(a)
a) CH3COOH, or b) H2CO3
 A 12.0 M solution of an acid that is
able to ionize completely in solution
would be termed:

( concentrated and strong )
Chapter 20 Review

According to the Bronsted-Lowry
theory, water will act as a base when
it:
( accepts a hydrogen ion )
Chapter 18 Review

What are the Bronsted-Lowry acids in
this equilibrium reaction?
CN1- + H2O  HCN + OH1( H2O and HCN )
Chapter 18 Review
Calculate the pH, and state if it is an
acid, base, or neutral solution:
 [ H+ ] = 1 x 10-9 ( pH = 9; it is basic )
 [ OH- ] = 1 x 10-10 ( pH = 4; it is acidic )
 [ OH- ] = 1 x 10-1 ( pH = 13; it is basic )
 [ H+ ] = 1 x 10-7 ( pH = 7; it is neutral )

Chapter 18 Review

Calculate the hydrogen-ion
concentration [ H+ ] for an aqueous
solution in which [ OH- ] is 1 x 10-12
mol/L. Is this solution acidic, basic,
or neutral?
( [ H+ ] = 1 x 10-2; it is an acidic solution )
Chapter 18 Review

Name the following acid, and then
write the expression for Ka. Assume
that only one hydrogen is ionized:
H2SO3
[ H+ ] x [ HSO3- ]
Ka =
[ H2SO3 ]
Sulfurous
acid
Chapter 18 Review

Name the following acid, and then
write the expression for Ka. Assume
that only one hydrogen is ionized:
HBr
Ka =
[ H+ ] x [ Br- ]
[ HBr ]
Hydrobromic
acid
Chapter 18 Review

Name the following acid, and then
write the expression for Ka. Assume
that only one hydrogen is ionized:
HNO3
[ H+ ] x [ NO3- ]
Ka =
[ HNO3 ]
Nitric
acid
Chapter 18 Review

Compare and contrast the properties
of acids and bases.
( Both acids and bases are electrolytes; they
cause indicators to change colors; and they
react with each other to form water and a
salt. Acids taste sour, bases taste bitter.
Bases feel slippery. Acids react with some
metals to produce hydrogen gas. )