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Chemistry 20
KEY
Chemistry 20
K.lVIogck
2016
.:msm.ogcksclassroom..coni
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Topic 8 Atoms and Molecules
Protons:
•
•
positively charged particles found inside the nucleus
determine protons by the atomic number
Electrons:
•
•
•
•
negatively charged electrons are found in a “cloud” around the nucleus (in specified
shells) which can be found be the period #
determine electrons by the atomic number
if the atom has a + charge, subtract that charge from the atomic number to find the
number of electrons
if the atom has a - charge, add that charge from the atomic number to find the number of
electrons
Neutrons:
•
•
neutral particles that make up mass within the nucleus
determine neutrons by subtracting the atomic number from the molar mass of an
element
Examples:
Element / Ion
# of Protons
# of Neutrons
# of Electrons
Ca
20
20
20
Al3+
13
14
10
Se
34
45
34
P2-
15
16
17
Electron Dot Diagrams
•
•
•
•
in chem. 20 we will NOT be dealing with the atom diagrams or Bohr model (only electron
dot diagrams)
electron dot diagrams show ONLY the electrons on the outer ring (VALENCE) of the
atom
remember to fill up the first lone orbitals and then start pairing
remember lone electrons interact and therefore will be next to each other
Examples:
Ionic Bonding:
Ionic bonds are composed of a metal element bonding with a nonmetal element
• The metal element contains a positive charge and is also known as a cation
• The nonmetal element has a negative charge and is also known as an anion
An ionic bond is created through the transfer of electrons. The metal element donates its extra
valence electrons to the non-metal to allow both elements to have a full valence. The molecule’s
composition can be determined through “drop and cross” method in which the charges of each
atom’s ions are transferred to the element it is bonding with as shown below:
Ca2+
P3-
à Ca3P2
Naming an ionic compound is the metal name followed by the nonmetal ending in “ide”. In the
example above the molecule is composed of calcium (the metal) and phosphorus (the nonmetal)
and therefore is named Calcium Phosphide.
Polyatomic ions are found at the top of your periodic table.
Examples:
For the listed compounds, draw the electron dot diagrams to demonstrate ionic bonding and
determine the molecular formula for each.
Molecular Bonding:
Molecular bonds are composed of 2 nonmetals (anions) creating a bond by sharing electrons.
When naming molecular compounds, prefixes are used to demonstrate number of atoms, and the
secondary atom ends with the suffix “ide”. It is important for you to memorize the prefixes
contained in the chart below.
Number
1
2
3
4
5
Prefix
Mono
Di
Tri
Tetra
penta
Number
6
7
8
9
10
Prefix
Hexa
Hepta
Octa
Nona
deca
Examples:
For the listed compounds, draw the electron dot diagrams to demonstrate molecular bonding and
determine the molecular formula or name for each.
Carbon dioxide
Methane
There are various molecular compounds that are expected knowledge for chemistry 20.
Please memorize the molecular formulas for the common molecular compounds in the chart below:
Name
Compound
Name
Compound
Ammonia
NH3 (g)
Glucose
C6H12O6 (s)
Sucrose
C12H22O11 (s)
Methane
CH4 (g)
Propane
C3H8 (g)
Acetic acid
CH3COOH (aq)
Water
H2O (l)
Ethanol
C2H5OH (l)
Methanol
CH3OH (l)
Ozone
O3 (g)
Hydrogen Peroxide
H2O2 (l)
Practice:
1. Complete the chart below
Name
Be
Ge
As3-
Ra1+
Xe
beryllium
Germanium
Arsenic ion
Radium ion
Xenon
Atomic #
4
32
33
88
54
Molar
Mass
9.01g/mol
72.64g/mo
74.92g/mol
226g/mol
131.29g/mol
Protons
Electrons
Neutrons
4
4
5
32
32
41
33
36
42
88
87
138
54
54
77
2. Draw electron dot diagrams for the following atoms or ions below:
3. In the space below, show the formula and the transfer of electrons demonstrating the ionic
bond using electron dot diagrams
4. In the space below, show the molecular formula and the sharing of electrons demonstrating the
molecular bond using electron dot diagrams
Topic 9 Periodic Trends and Ionization
An electrostatic force is a force existing as a result of the attraction or repulsion between 2
charged particles. There are 4 important bonding relationships to remember:
1. Opposite charges attract to each other.
2. Like charges repel each other.
3. The greater the distance between 2 charged particles, the
smaller the attractive (or repulsive) force
4. The greater the charge, the greater the force of attraction
(or repulsion) between them
Across a Period
Periods travel left to right, horizontally across a periodic table. Each period on the periodic table
represents a different layer of electrons or a different “electron shell”. The 1st shell has 2
electrons (therefore the 1st period has 2 elements) the 2nd and 3rd shells have 8 electrons
(therefore the 2nd and 3rd period have 8 elements).
Left to right the atomic number (# of protons in the nucleus) increases and the positive charge on
the nucleus increases. This increases the number of electrons around a nucleus. All the electrons
in a given shell can be assumed to have the same average distance from the nucleus.
Valence Electrons and Valence
Valence electrons are electrons in open shells (not full shells). Valence electrons are reactable
electrons. Noble gases have no valence electrons. From this we can determine that valence
electrons make an element reactive.
In an atom, each individual orbital holds up to 2 electrons… (Totaling 8 valence electrons). Since
electrons repel each other, each electron added goes into a vacant orbital and only pairs if there
are no empty orbitals remaining. Valence can be noted as the number of electrons, which are
normally available for bonding or the number of unpaired electrons in an atom. Valence is also
known as combining capacity or bonding capacity.
Ionization Energy
In order to form a positive ion, an electron must be removed from a neutral atom. For example to
create Li+ you must add energy to pull off one electron from Li(s). Ionization energy is the energy
required to remove an e- from a neutral atom. The e- removed is always a valence unless the atom
has a closed shell.
Using what you know about periodic trends, state the relationship of ionization energy on the
periodic table below. (Please give a good explanation).
Going across a period, the number of protons increases, making
the nucleus have a higher charge and be more attractive. This,
accompanied by minimal shielding will make fluorine the most
attractive and thus require the most energy to remove the
electrons.
Going down the group, the number of orbitals increases, causing
the shielding to increase and the distance between the nucleus
and the out electrons large. This allows the electrons to be easily
removed from the outer shell, with minimal energy input
Practice Problems:
1. What happens to the atomic radius (the distance between the nucleus and the outermost
electrons) as you proceed down a group (top to bottom)? Explain this trend.
Atomic radius Increases down a group as more shells are
added each period
2. What happens to the charge of the nucleus from left to right across a period? What effect
would you expect this change in charge to have on the average distance between the
nucleus and electrons? Explain your answer completely.
Across a period more protons are in the nucleus, creating
greater charge and therefore greater attractive forces
3. Which family of elements appears to possess only closed shells?
Noble Gases
4. Circle which of the following have an open shell and place a C next to those that have
closed shells?
a. Cl - open
b. Ne - closed
c.
Mg - open
d. Si - open
e.
Na+ - closed
f.
Cl- - closed
g.
O- - open
h. Ca2+ - closed
i.
I - open
j.
Al+ - open
5. Fill in the number of valence electrons corresponding with each atom
Atom
F
Ne
Na
Ne+
# Of valence e-
7
8 or 0
1
7
Atom
Pb
Pb2+
SS2-
# Of valence e-
2
8 or 0
7
8 or 0
6. In the space below, draw the electron dot diagrams for each of the elements (or ions) in
the chart from question 5.
F
Pb
Ne
Pb2+
Na
S1-
Ne1+
S2-
7. What happens to the electrostatic attraction of the nucleus to an electron in the outmost
shell going down a group? Explain this trend.
Going down a group, the electron is less attracted to the
nucleus, due to the increased distance and increased
shielding within the atom.
What happens to the ionization energy going down a group?
Decreased as it is easier to remove an electron
8. What happens to the nuclear charge going across a period and how does that effect the
attraction to the electrons in the outer shell? Explain this trend.
Nuclear charge increases (more protons = more positive
charge) causing the nucleus to be more attractive to the
valence electrons.
9. What happens to the ionization energy going across a period?
Increased as the electrons are more attracted to the
nucleus
10. Which member in each of the following pairs should have greater ionization energy?
Circle the atom with the greater ionization energy.
a. Br or Cl
b. Al or Cl
c. Ne or Xe
d. Mg or Ba
e. F or Ne
f. Rb or I
Topic 10 Chemical Reactions
It is expected out of science 10 that every student is able to identify the type of chemical reaction,
as well as balance chemical reactions. This is a very important skill set for chemistry 20. The
chart below describes the 6 types of chemical reactions.
Determine a NON-CHEMISTRY way of helping you remember these reactions and draw them in
the space provided!
Formation
•
•
2 elements combine to form 1
compound
2Na + Cl2 à NaCl
Decomposition
•
•
Single Replacement
•
•
an element reacts with a compound,
creating a different compound and
an element
Zn + Cu(NO3)2 à Zn(NO3)2 + Cu
Hydrocarbon Combustion
•
•
a hydrocarbon compound burns in
the presence of oxygen to produce
carbon dioxide and water
CH4 + 2O2 à CO2(g) + 2H2O(l)
1 compound breaks down into its
elements
2HgO à 2Hg + O2
Double Replacement
•
•
2 compounds react to form 2
different compounds
AgCl + MgF2 à AgF + MgCl2
Acid Base Neutralization
•
•
an acid (H) reacts with a base (OH)
to produce water and an ionic
compound
NaOH + H2SO4 à 2H2O + Na2SO4
Solubility:
Solubility determines whether a compound will be insoluble and therefore produce a precipitate
(solid) after a reaction, or be soluble and therefore remain aqueous.
To determine solubility, use a solubility chart. Find the anion (negative ion) across the top of the
solubility chart. Find the attached cation (positive ion) in the boxes below.
• If the cation is found in the top section, the compound will be soluble and therefore remain
in the aqueous state.
• If the cation is found in the lower section, the compound is insoluble and therefore will
become a solid precipitate.
The solubility chart we will refer to in chemistry 20 is found on your periodic table. Using the
solubility chart, determine whether the following examples are solid or aqueous in solution.
Examples:
1. Using the solubility chart, determine whether the following examples are solid or aqueous in
solution.
a. MgF2 (s)
e. CuCl (s)
b. K2CO3 (aq)
f. CuCl2 (aq)
c. Ca(OH)2 (s)
g. PbSO4 (s)
d. NaCl (aq)
h. NH4IO3 (aq)
Balancing:
Often times students find balancing equations the hardest part in chemistry! If you follow this
step by step guide to balancing, hopefully it should get a little easier!
1. Start with an atom involved in one species on each side of the equation. Assign coefficients
that balance this species. The first placement of coefficients should be the ONLY TIME you should
have to write in 2 coefficients. Metal atoms often dictate a reaction, so should be balanced first. If
there is no metal present look for anything that is not H or O. (they may eventually balance
themselves)
2. Select one of the atoms in the same molecule you have just balanced that is present once more
(without a coefficient) and add a single balancing coefficient.
3. Continue to find an atom that occurs once without a coefficient in front of it and balance it.
4. Double check that all atoms are balanced (start with the last atom to balance – usually by
default).
HINT: Until you have put a number in front of a species assume you have zero atoms (or
molecules) of this species. Treat a blank as a zero.
HINT: Try to balance entire groups (i.e. PO4, NO3, SO4, NH3) if possible!
HINT: If a fraction occurs during balancing, multiply the entire equation by the whole number
(i.e. 2) to eliminate the fraction
HINT: Remember DIATOMIC ELEMENTS occur as pairs!
Example:
Balance the following equations.
1(NH4)3PO4 + 3NaOH à 1Na3PO4 + 3NH3 + 3H2O
Reactants
N=
H=
P=
O=
Na =
Products
N=
H=
P=
O=
Na =
__4__C19H17NO3 + _87___O2 à __76__CO2 + __34__H2O + __2__N2
_1_Cr2(SO4)3 + _5_KI + _1_KIO3 + _3_H2O à _2_Cr(OH)3 + _3_K2SO4 + _3_I2
__4__MoCl3 + __1__O2 + __4__AgCl à __4__MoCl4 + __2__Ag2O
Practice:
Balance the following reactions. Determine what type of reaction is occurring and the state for
each compound.
1. __2__NaBr(aq)+ __1__Ca(OH)2(s) à __1__CaBr2(aq) + __2__NaOH(aq)
Type of reaction: Double replacement
2. __2__NH3(aq) + __1__H2SO4(aq) à __1__(NH4)2SO4(aq)
Type of reaction: Formation
3. __4__C5H9O(l) + _27___O2(g)à __20__CO2( g) + __18__H2O(g)
Type of reaction: Combustion
4. __3__Pb(s) + __2__ H3PO4(aq) à __3__H2(g) + __1__Pb3(PO4)2(s)
Type of reaction: Single replacement
5. __1__Li3N(aq) + __3__NH4NO3(aq) à __3__LiNO3(aq) + __1__(NH4)3N(aq)
Type of reaction: Double replacement
6. __3__HBr(aq) + __1__Al(OH)3(s) à __3__H2O(g) + __1__AlBr3(aq)
Type of reaction: Acid base neutralization
7. What is the main difference between a double replacement reaction and an acid base reaction?
Acid base is a specific type of double replacement reaction that
always produces water and an ionic substance (salt)
8. Combustion reactions always result in the formation of water. What other types of chemical
reactions may result in the formation of water. Give examples of these reactions.
Acid base neutralization
Formation
CH3COOH + NaOH à H2O + NaCH3COO
2H2+ O2 à 2H2O
Complete the following chemical equations, balance and determine the type of reaction when
complete.
9. __1___Fe(s) + ___1__H2SO4(aq) à 1 FeSO4(aq) + 1 H2(g)
single replacement
10. __2___C2H6(l) + __7___O2(g) à 4 CO2(g) + 6 H2O (l)
combustion
11. __3___KOH(aq) + ___1__H3PO4(aq) à 1 K3PO4(aq) + 3 H2O(l)
12. ___1__SnO2 (l) + __2___H2(g) à 1 Sn(s) + 2 H2O (l)
acid base neutral.
single replacement
13. __2___KNO3(aq) + __1___H2CO3(aq) à 1 K2CO3(aq) + 2 HNO3(aq)
double replacement
14. __1___B2Br6(l) + __6___HNO3(aq) à 2B(NO3)3(aq) + 6 HBr(aq)
double replacement
15. __2__ SeCl6(aq) + __1__ O2 (g) à 2 SeO(aq) + 6Cl2(g)
single replacement