Download Ionic and Covalent Compounds: Naming, Formulas, Properties 1

Ionic and Covalent Compounds: Naming, Formulas, Properties
WWW.CHEMISTRYMATTERS.SHUTTERFLY.COM
PASSWORD: chem
1 Write the formula or name for these compounds.
(A) CS 2
(F) carbon tetrabromide
(B) Cl 2 O 7
(G) diphosphorus trioxide
(C) chromium(III) nitrite
(H) sodium perchlorate
(D) Mg 3 (PO 4 ) 2
(I) Pb(C 2 H 3 O 2 ) 2
(E) LiF
(J) tetraiodine nonoxide
2 You are given the chemical formula, Cu 2 (SO 4 ) 3 . Answer the following questions regarding this formula.
(A) What is the charge of the cation?
(B) What is the name of the compound?
(C) List the ions and tell how many of each ion are in the formula.
(D) How many atoms of each element are there?
(E) What is a formula unit?
(F) If you have 1.5 moles of Cu 2 (SO 4 ) 3 , how many formula units do you have? (Show work!)
(G) What are three generic physical properties of this ionic compound?
Name
STUDY GUIDE
Date
Class
6.2 Types of Reactions
CHAPTER
Use with text pages 202 – 209
6
Choose the correct symbol listed below to describe each of the chemical
reactions listed in questions 1 – 15. Also write a balanced chemical equation
for each reaction.
S = synthesis
D = decomposition
C = combustion
1.
SD = single displacement
DD = double displacement
zinc(s) 1 silver nitrate(aq) ˇ silver(s) 1 zinc nitrate(aq)
balanced equation:
2.
iron (III) hydroxide(s) ˇ iron (III) oxide(s) 1 water(g)
balanced equation:
3.
ammonia(g) 1 sulfuric acid(aq) ˇ ammonium sulfate(aq)
Copyright © Glencoe / McGraw - Hill, a division of The McGraw - Hill Companies, Inc.
balanced equation:
4.
magnesium(s) 1 nitric acid(aq) ˇ magnesium nitrate(aq) 1 hydrogen(g)
balanced equation:
5.
calcium nitrate(s) 1 hydrochloric acid(aq) ˇ nitric acid(aq) 1
calcium chloride(aq)
balanced equation:
6.
Na(s) 1 H2O(l) ˇ NaOH(aq) 1 H2(g)
balanced equation:
7.
CO(g) 1 O2(g) ˇ CO2(g)
balanced equation:
8.
FeS(s) 1 HCl(aq) ˇ FeCl2(aq) 1 H2S(g)
balanced equation:
Chemistry: Concepts and Applications
Study Guide, Chapter 6
23
Name
Date
9.
Class
NaNO3(s) ˇ NaNO2(s) 1 O2(g)
balanced equation:
10.
CH4(g) 1 O2(g) ˇ CO2(g) 1 H2O(g)
balanced equation:
11.
Fe(s) 1 CuNO3(aq) ˇ Cu(s) 1 Fe(NO3)2(aq)
balanced equation:
12.
KI(aq) 1 Cl2(g) ˇ KCl(aq) 1 I2(aq)
balanced equation:
13.
Al(s) 1 S(s) ˇ Al2S3(s)
balanced equation:
14.
KClO3(s) ˇ KCl(s) 1 O2(g)
balanced equation:
C4H10(g) 1 O2(g) ˇ CO2(g) 1 H2O(g)
Copyright © Glencoe / McGraw - Hill, a division of The McGraw - Hill Companies, Inc.
15.
balanced equation:
24
Chemistry: Concepts and Applications
Study Guide, Chapter 6
WWW.CHEMISTRYMATTERS.SHUTTERFLY.COM
PASSWORD: chem
P1V1 / T1 = P2V2 / T2
R = 0.0821 atm L mol-1 K-1 PV = nRT
n = m / MM 1 atm = 101.3 kPa = 760 mmHg
K
KEY: P = pressure; V = volume; T = temperature; R = gas constant; n = number of moles; MM = molar mass; m =mass
°C + 273 =
All units and work must accompany problem in order to receive full credit!!!
1. The pressure is doubled on 1000 mL of gas at the same time the temperature is changed from –86 °C to 86 °C.
Solve for the final volume in liters. Show ALL steps to the solution including before and after substitution into
formulas.
2. A 50.0 L hot air balloon was filled with 4035 g of H 2 (hydrogen gas) on a hot balmy day (30.0◦C). After 1 hour, 15.0%
(or 0.150 of the total) of the H 2 had escaped from the balloon. How many molecules of hydrogen gas remained? If
the final volume of the balloon is 30.0 L, what is the pressure (in kPa) after it deflates? Show ALL steps to solution
including any conversions (including molar mass).
3. Ammonia gas is reacted with oxygen gas to produce water vapor and nitrogen monoxide gas. The UNBALANCED
reaction is: ____NH3(g) + ____O2(g)  ____NO(g) + ____H2O(g). If 68.0 grams of ammonia gas are placed in a
50.0 L container at 40.0 ºC, what is the pressure of nitrogen monoxide produced? What is the total pressure inside
the container after the reaction takes place? Assume the container was initially filled with 68.0 grams of ammonia gas
and exactly enough oxygen to react with it. Show ALL steps to the solution.
4. If 600. ml of hydrogen gas is produced through a replacement reaction involving solid iron and sulfuric acid at STP, how
many grams of Iron (III) sulfate are produced? Balanced Equation: 2Fe + 3H 2 SO 4  3H 2 + Fe 2 (SO 4 ) 3
5. If 14.7 grams of sulfuric acid (H 2 SO 4 ) were used in the above reaction, how many liters of hydrogen would be collected
at STP? Balanced Equation: 2Fe + 3H 2 SO 4  3H 2 + Fe 2 (SO 4 ) 3
6. When 950. mL of the hydrogen gas was collected over water at a temperature of 25 Celsius and at a barometric
pressure of 765 mmHg, how many grams of sulfuric acid were needed? Balanced Equation: 2Fe + 3H 2 SO 4  3H 2 +
Fe 2 (SO 4 ) 3
7. If using the above equation 95.6 grams of solid Fe were used, what volume of hydrogen gas would be produced at 1.05
atm and 22 Celsius? Balanced Equation: Fe + H 2 SO 4  H 2 + Fe 2 (SO 4 ) 3
Moles, Stoichiometry, Gases REVIEW
WWW.CHEMISTRYMATTERS.SHUTTERFLY.COM
PASSWORD: chem
Useful Equations: PV = nRT (R = 0.0821 atm L mol-1 K-1) P 1 V 1 /T 1 = P 2 V 2 /T 2
n = m/MM
T K = T C +273
Where P is the pressure, V is the volume, n is the number of moles, T is the temperature, R is
the universal gas constant, m is the mass given, and MM is the molar mass.
Yo u ar e a NASCAR p it cr ew m em b er . Yo u r p it ’s car is lead in g t h e r ace w it h 20 lap s t o g o . Yo u ar e p r ayin g
t h at t h e car h as en o u g h g as t o f in ish t h e r ace. Yo u ad d an alt er n at iv e f u el, iso p en t an e, t o t h e f u el t an k.
The
unbalanced
reaction
is
as
follows:
C 5 H 12 (g)
+
For each of the parts below, be sure to show ALL work including units.
(example: moles of O 2 = blah. Moles of CO 2 = blah, etc.)
O 2 (g)

CO 2 (g)
+
H 2 O(g)
Also, label portions of your answer
a) Balance the reaction above.
___ C 5 H 12 (g) + ___ O 2 (g)  ___CO 2 (g) + ___H 2 O(g) [8 points]
b) What is the mole ratio of oxygen gas to carbon dioxide gas? ___mol O 2 :___mol CO 2 [10 points]
c) If one mole of oxygen gas reacts, how many moles of carbon dioxide gas are produced? [8 points]
d) What are the molar masses of isopentane, C 5 H 12 , AND carbon dioxide, CO 2 ? [8 points]
e) The volume of carbon dioxide produced during the race is 63.2 liters assuming conditions are at STP. What is
the mass in grams of isopentane, C 5 H 12 , AND the mass in grams of oxygen gas, O 2 , necessary to produce this
volume of carbon dioxide? Show ALL work and label ALL steps. [18 points]
f) The volume of carbon dioxide produced during the race is 63.2 liters. What is the mass in grams of isopentane,
C 5 H 12 , AND the mass in grams of oxygen gas, O 2 , necessary to produce this volume of carbon dioxide? The
atmospheric pressure is 695 mm Hg and the temperature is 40.0 degrees Celsius. Show ALL work and label
ALL steps. [26 points]
g) The volume of carbon dioxide produced during the race is 63.2 liters. If conditions were initially at STP, what
would the volume be if the pressure were at 695 mm Hg and the temperature was 40.0 degrees Celsius?
Show ALL work and label ALL steps. [10 points]
h) If 216.0 grams of isopentane, C 5 H 12 , reacted with 128.0 grams of oxygen gas, O 2 , what is the limiting
reactant? How much (in grams) of the excess reactant is unused? Show ALL work and label ALL steps. [12
points]
i)
BONUS: If 216.0 grams of isopentane reacted with 128.0 grams of oxygen gas, how much water (molar mass:
18.0 g/mol) would be produced? [10 points]
Determining Empirical &
Molecular Formulas
Calculating Molar Mass
Calculate the formula mass of
magnesium carbonate, MgCO3.
The total
atomic
mass of the
compound
in grams
24.31 g + 12.01 g + 3(16.00 g) = 84.32 g
Calculating Percentage Composition
Calculate the percentage composition of magnesium
carbonate, MgCO3.
From previous slide:
24.31 g + 12.01 g + 3(16.00 g) = 84.32 g
 24.31 
Mg
= 
= 28.83%
 • 100
 84.32 
 12.01 
=
C 
= 14.24%
 • 100
 84.32 
 48.00 
=
O 
= 56.93%
 • 100
 84.32 
100.00%
Formulas
Empirical formula: the lowest whole number
ratio of atoms in a compound.
Molecular formula: the true number of
atoms of each element in the formula of a
compound.
molecular formula = C6H6
 empirical formula = CH
Formulas (continued)
The empirical and molecular formula for
ionic compounds are ALWAYS the same.
Examples:
NaCl
MgCl2
Al2(SO4)3
K2CO3
Formulas (continued)
Formulas for molecular compounds MIGHT
be empirical (lowest whole number ratio).
Molecular:
H2O
C6H12O6
C12H22O11
Empirical:
H2O
CH2O
C12H22O11
Empirical Formula Example
Adipic acid contains 49.32% C, 43.84% O, and
6.85% H by mass. What is the empirical formula
of adipic acid?
( 49.32 g C )(1 mol C ) = 4.107 mol C
(12.01 g C ) 2.74 mol
( 6.85g H )(1 mol H ) = 6.78 mol H
(1.01 g H ) 2.74 mol
( 43.84 g O )(1 mol O ) = 2.74 mol O
(16.00 g O ) 2.74 mol
= 1.50
Not
whole
#s
= 2.47
= 1
smallest value
Empirical Formula Example
Multiply each number by an integer to obtain all
whole numbers.
Carbon: 1.50
x 2
3
Hydrogen: 2.47
x 2
5
Oxygen: 1.00
x 2
2
Empirical formula: C3H5O2
Calculation the Empirical Formula
Steps:
1. Convert the mass of each element
to moles.
2. Divide each mole value by the
smallest of the mole values.
3. Multiply each number by an
integer to obtain all whole
numbers (if required).
Calculating the Molecular Formula
The empirical formula for adipic acid is
C3H5O2. The molecular mass of adipic acid is
146 g/mol.
What is the molecular formula of adipic acid?
1. Calculate the mass of the empirical formula
C3H5O2
C
H
O
3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g
The empirical formula for adipic acid is
C3H5O2. The molecular mass of adipic acid is
146 g/mol.
What is the molecular formula of adipic acid?
3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g
2. Divide the molecular mass by the
empirical mass.
146
=2
73
The empirical formula for adipic acid is
C3H5O2. The molecular mass of adipic acid is
146 g/mol.
What is the molecular formula of adipic acid?
3(12.01 g) + 5(1.01) + 2(16.00) = 73.08 g
146
=2
73
3. Multiply the empirical formula by
this number to get the molecular
formula.
(C3H5O2) x 2
=
C6H10O4
AP Chemistry Summer Assignment for 2013-14 School Year
1. MEMORIZE the polyatomic ions and charges of most commonly used transition metals
(see table below) and memorize the few formulas for nonmetals.
POLYATOMIC IONS
CATIONS (positive ions)
+1
+2
GROUP IA ELEMENTS (Alkali
+
+
+
Metals): Li , Na , K , etc.
+4
+3
GROUP IIA
IONS OF CERTAIN GROUP
IIIA ELEMENTS, including
3+
3+
Al & B
ELEMENTS (Alkaline
2+
2+
Earth Metals): Mg , Ca , etc.
+
SILVER, Ag
LEAD(IV), Pb4+
MOST TRANSITION METALS FORM
IONS WITH +2 CHARGES
AMMONIUM, NH 4
+
TIN(IV), Sn4+
IRON(III), Fe
3+
LEAD(II), Pb2+
TIN(II), Sn
2+
ANIONS (negative ions)
-1
-2
GROUP VIIA ELEMENTS (the
- Halogens): F , Cℓ , Br , I
ACETATE, CH 3 COO
2-
BICARBONATE, HCO 3
(hydrogen carbonate)
-
BISULPHATE, HSO 4
(hydrogen sulphate)
3-
CARBONATE, CO 3
2-
CHROMATE, CrO 4
2-
OXALATE, (COO) 2
PEROXIDE, O 2
-
HYDROXIDE, OH
NITRATE, NO 3
SULPHITE, SO 3
-
3-
2-
2-
2-
2-
THIOSULPHATE, S 2 O 3
-
2-
NONMETALS
-
PERCHLORATE, CℓO 4
-
You should know that the following nonmetals
are diatomic: H2, N2, O2, F2, Cl2, Br2 & I2.
-
PERMANGANATE, MnO 4
THIOCYANATE, SCN
PHOSPHATE, PO 4
2-
SULPHATE, SO 4
-
CYANATE, OCN
GROUP VA
ELEMENTS, including N
DICHROMATE, Cr 2 O 7
CYANIDE, CN
2-
GROUP VIA ELEMENTS: O , S , etc.
-
CHLORATE, CℓO 3
-4
-3
-
Phosphorus exists as P4; phosphorus oxide can
exist either as P4O10 (most likely) or as P4O6.
Sulphur normally exists as S8 molecules.
2. MEMORIZE the gas information below:
Volume is always measured in litres (L), millilitres (mL) or cubic centimetres (cm3). 1 L = 1000 mL = 1000 cm3
DENSITY OF A GAS = MOLAR MASS OF THE GAS (g/mol)
(g/L)
MOLAR VOLUME OF THE GAS (L/mol)
The five principal assumptions of the kinetic molecular theory of gases are as follows:

Gases consist of molecules whose volumes are negligible compared with the volume occupied by the
gas.

Since the molecules of a gas are far apart, the forces of attraction between them are negligible.

The molecules of a gas are in continual, random, and rapid motion.

The average kinetic energy of gas molecules depends only on the gas temperature, and can be
expressed by E K α T.

Gas molecules collide with each other and with the walls of their container, but they do so without loss
of energy (The collisions are said, by scientists, to be "perfectly elastic").
Real Gases versus Ideal Gases
The Gas Laws work for “Ideal Gases”. However, there is no such thing as an Ideal Gas!
•
Real Gases deviate most from ideal behaviour at high pressures.
•
Real Gases deviate most from ideal behaviour at low temperatures.
•
At a given temperature and pressure, the greater the intermolecular forces, the greater will be the
deviation from ideal behaviour.
At high pressures and low temperatures, the greater the size of the gas molecules, the greater will be the
deviation from ideal behaviour.
•
3. KNOW the electron configuration information below:
When writing electron configurations, the following “rules” govern the placement of each electron in the
electron cloud of the atom:
The Aufbau Principle: This simply states that the lowest energy level orbitals are filled first.
The Pauli Exclusion Principle: This states that no two electrons in an atom can have the same set of four
quantum numbers.
Hund’s Rule: This states that the most stable arrangement of electrons is that with the maximum number of
unpaired electrons, all with the same spin direction.
subshell
s
p
d
f
maximum
number of
electrons
2
6
10
14
Filled subshells, particularly in the valence shell, lead to stable, unreactive, elements. Additionally, there is
stability associated with half-filled valence subshells.
The Noble Gases are, of course, the most stable of all the elements.
Nitrogen, phosphorus and arsenic have unusually stable properties due to the fact that their valence shells
consists of filled s orbitals and half-filled p orbitals.
4. KNOW periodic trends below:
Periodic Trends: Sizes of Atoms, Ionization Energies, Electronegativity
1) Within each column (group), the atomic radius tends to increase as we proceed from
top to bottom.
2) Within each row (period), the atomic radius tends to decrease as we move from left
to right.
The ionization energy of an atom or ion is the minimum energy required to remove an electron from the
ground state of the isolated gaseous atom or ion. The first ionization energy, I 1 , is the energy needed to remove
the first electron from a neutral atom. For example, the first ionization energy for the sodium atom is the energy
required for the process, Na(g)  Na+(g) + eThe second ionization energy, I 2 , is the energy needed to remove the second electron, and so forth, for
successive removals of additional electrons.
1. Within each group, ionization energy generally decreases with increasing atomic
number as you proceed down the group.
2. Within each row, ionization energy generally increases with increasing atomic
number as you proceed from left to right. There are slight irregularities in this
trend, however.
Electronegativity is the ability of an atom in a molecule to attract electrons to itself.
Electronegativities generally decrease as you proceed down a group and increase as you proceed from left to
right across a period. Fluorine is the most electronegative element.
Ionic Bonding & Ions
An ionic bond is formed when one electron, or more, is/are transferred from one atom to another.
Positive ions are referred to as cations; negative ions are referred to as anions.
Lattice energy is governed by the formula,
E∝
Q1Q2
d
(where E is the lattice energy, Q 1 and Q 2 are the charges of the ions and d is the distance between them).
The most common ionic charge of the fourth row transition elements, from scandium to zinc, is 2+ (e.g. Ti2+ &
Fe2+). When such ions are formed, the transition metal atom loses its two 4s electrons (3d electrons are not
lost). (In fact, whenever a positive ion is formed from an atom, electrons are always lost first from the subshell
having the largest value of n). Thus, in forming ions, transition metals lose the valence-shell s electrons first,
then as many d electrons as are required to reach the charge of the ion.
Relative Sizes of Ions
For the Representative (s-block and p-block) Elements that form positive ions (cations), the radius of
the positive ion will always be smaller than the radius of the corresponding atom. This is due primarily
to the fact that when these elements form ions the outermost shell (highest value of n) is lost in its
entirety.
For the Representative Elements that form negative ions (anions), the radius of the negative ion will
always be larger than the radius of the corresponding atom.
For all of the Representative Elements, as you go down a group the radii of ions of equal charge
increase. This is due primarily to the fact that as you go down the group the outermost electrons have
a larger value of n.
Structures for Covalent Molecules

Ensure that each atom ends up with the “correct” number of valence electrons.
Most elements end up with eight valence electrons (the “octet rule”).
Hydrogen ends up with two valence electrons.
Boron and beryllium usually end up with fewer than eight valence electrons.
Some elements from periods 3, and higher, end up with more than eight valence
electrons. These elements include sulphur, phosphorus, arsenic, selenium and
xenon. They are said to form “expanded octets”.