Download Unit 2 Exam Review 2

PreIB/AP Chemistry ~ Atomic Structure and the Periodic Table Unit Review
Atomic structure
A great number of scientists have contributed to the development of the atomic theory. Our currently held theory
is open to change as new technologies improve our abilities to make observations and experiment. Identify the
scientists who made the following contributions:
______________________= all substances are composed of atoms which cannot be created, divided or destroyed
______________________=organized the known elements in to a Periodic Table
______________________= Plum Pudding Model = positively charged sphere with embedded electrons
______________________=discovered positively charged nucleus by shooting alpha particles at gold foil
______________________= proposed that electrons have to be in energy levels around the nucleus, energy is
released or gained when electrons jump between levels
Currently, our theory of the structure of the atom holds that atoms have a nucleus that is positively charged and
contains protons and neutrons. Electrons move freely around the nucleus but statistically can be found most often
in spaces called orbitals. Only two electrons can ‘fit’ in an orbital; groups of orbitals are organized into energy levels
and sub-levels.
The Periodic Table
The periodic table is arranged by increasing atomic number. The columns (up and down) are called groups (or
families). The rows (sideways) are called periods.
Elements in the same group have similar properties because they have the same number of valence electrons. For
example, Group 1 elements have 1 valence electron. Group 13/3A elements have 3 valence electrons. Elements in
the same period are filling the same valence energy level. For example, elements in period 4 all have valence
electrons in the 4th energy level.
Some groups have special family names:
Group 1 = Alkali metals – 1 valence electron - the most reactive of all metals
Group 2 = Alkaline earth metals – 2 valence electrons
Groups 3-12 = Transition metals – these elements all have their last electron in the d block
Group 17 = Halogens = 7 valence electrons - the most reactive of non-metals
Group 18 = Noble Gases = 8 valence electrons – the least reactive of all elements due to their full octet of valence
electrons; also called inert gases due to their inactivity.
An ion is an atom that has gained or lost electrons; due to this difference, ions have charges. Ions with 1 extra
electron will have a charge of -1 (because electrons carry a negative charge). Ions with 1 fewer electron will have a
charge of +1. Oxidation number is the charge an element is likely to have if it’s an ion. The oxidation number
can also be thought of as the number of electrons the element needs to gain or lose to have a full valence energy
level. For example, the oxidation number of oxygen is -2 because oxygen has 6 valence electrons and needs 2 more
to make 8 (octet = full energy level). Sodium has an oxidation number of +1 because it’s easier to lose one valence
electron than it is to get 7 more electrons to make 8.
Complete the data table:
Group #
1
2
17
18
Questions:
Family Name
# Valence Electrons
Oxidation #
1.
2.
3.
4.
5.
6.
In what family (family name, not group number) is calcium? ___________________________
In what family is krypton? ________________________________
In what family is lithium? ______________________________
How many valence electrons does chlorine have? ____________
What is the oxidation number of nitrogen? ___________
Which of the following groups of elements have similar properties?
a. N, Ne, Ni
b. Sr, Ba, Be
c. C, N, O
d. Ru, Rh, Ir
7. Why do elements in the same group have similar chemical and physical properties?
__________________________________________________________________________________
Isotopes are elements (and so will have the same number of protons) with different numbers of neutrons and
therefore different masses.
Three isotopes of lithium:
Ions are atoms with electrical charges due to different numbers of protons and electrons
Charge = (number of protons) – (number of electrons)
cation = positive charge (more protons than electrons)
anion = negative charge (more electrons than protons)
neutral atom = zero charge (protons equal electrons)
Molar mass
Molar mass is found with the atomic mass numbers on the periodic table. Atomic mass or molecular mass is
labeled in amu’s (atomic mass units). Molar mass is labeled in grams/mole and it is the mass of one mole of the
substance.
Sample questions:
1.
2.
3.
4.
5.
What is the atomic mass of sodium?
What is the molar mass of sodium?
What is the molar mass of chlorine?
What is the molar mass of sodium chloride, NaCl?
What is the molar mass of ammonium nitrate, NH4NO3?
Average Atomic Mass is a calculation that shows the average mass of the isotopes of an element taking into
account the abundance of each isotope. To calculate average atomic mass:
1 – Change the percent abundance to a decimal (divide by 100)
2 – Multiply each abundance (decimal) by the mass of that isotope
3 – Add the products from step 2
Three isotopes of an element are shown below. Calculate the average atomic mass. Draw a box around your final
answer.
Isotope
X-29
X-31
X-32
Abundance
2.0%
11%
87%
Mass (amu)
29.2
31.4
32.5
Percent composition is a calculation that determines the parts of a compound that each element comprises. To
calculate percent composition, use the formula below:
𝑝𝑎𝑟𝑡
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡
Percent Composition = 𝑤ℎ𝑜𝑙𝑒 × 100 = 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑡𝑜𝑡𝑎𝑙 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑 × 100
Find the percent compositions (both carbon and oxygen) of CO2
Electron Configurations
Electron configurations show us an atom’s arrangement of electrons. Knowing how electrons are arranged helps us
predict characteristics about that atom. There are three rules for determining the arrangement of electrons around
the nucleus of an atom: 1) Electrons must be placed in orbitals with the lowest energy first (Aufbau). 2) Two
electrons can be placed in each orbital but must have opposite spins (Pauli). 3) If there are orbitals of equal energy,
one electron is placed in each before two can be placed in any (Hund). There are different numbers of orbitals in
each sub-level and energy level.
Energy Levels
Principle Quantum
Number (period #)
1
Sub-Levels
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
Number of Orbitals
1
1
3
1
3
5
1
3
5
7
Total Electrons
2
2
6
2
6
10
2
6
10
14
2
3
4
We can show electron configuration with arrows. Complete the data table below; represent electrons with arrows.
Element
Oxygen
Sodium
Argon
1s
2s
2p
3s
3p
4s
We can also show the numbers of electrons in each level with exponents over the orbital labels.
For example, 1s22s22p3 shows us that there are 2 electrons in the 1s orbital, 2 electrons in the 2s orbital, and 3
electrons in the 2p orbital. That is a total of 7 electrons (2+2+3), which means the element is nitrogen (atomic
number 7).
We can also use the electron configuration to determine the number of valence electrons. The number of valence
electrons is always equal to the sum of the electrons in the s and p orbitals in the last energy level. For example,
1s22s22p3 has 5 valence electrons, because in the valence (highest) energy level (2) there are 2+3 electrons. (For our
purposes, D and F orbitals do NOT count as valence electrons.)
We can also identify the group an atom is in by its electron configuration. 1s22s22p3 has 5 valence electrons, so it’s
in Group 5A or 15. Perhaps even easier, we can use the blocks of the periodic table to predict the group. The last
electron of 1s22s22p3 is in the 2p3 space, which is the 2nd period, p block, 3rd column, which is group 15 or 5A.
Elements in the same group will have the same configuration for their last electrons. For example, elements in
group 1 will all end in s1. Lithium is 1s22s1. Potassium is 1s22s22p63s23p64s1.
Sample questions:
1.
2.
3.
4.
5.
What element has the electron configuration 1s22s22p1? ________________________
What element has the electron configuration 1s22s22p63s23p6? ________________________
How many valence electrons does 1s22s22p1 have? ________
In what family (family name, not group #) is 1s22s22p63s23p5? ____________________
An element has an electron configuration of 1s22s22p63s23p1. Which of the following will be in the same
group as this element?
a. 1s22s22p63s23p3
b. 1s22s22p63s1
c. 1s22s22p63s23p63d104s24p1
Lewis Dot Structures
Lewis Dot Structures show us how many
valence electrons an element has. Each valence
electron is represented by a dot. We can use
these structures to predict how atoms will bond
with each other.
Elements in group 1 have one valence electron,
so their Lewis Dot Structures will each have
one dot. Group 2 will have 2 dots. Group
18/8A will have 8.
Sample questions:
1. Draw the Lewis Dot Structure of Boron:
2. How many valence electrons would an element with the following Lewis Dot Structure have? _____
X
3. Which Lewis Dot Structure represents a halogen atom in the ground state?
a.
X
b.
X
c.
X
d.
X
Periodic Trends
Periodic Trends are the patterns that we find in the periodic table. This is also known as periodicity. There are
three important trends we focus on this year:
-
-
Atomic radius is the size of an atom. Atoms on the left side of the periodic table tend to be bigger than atoms
on the right because as we add another proton to each element going from left to right, the magnetic attraction
between the positive protons in the nucleus and the electrons in the energy levels becomes stronger and the
protons pull the electrons closer to the nucleus. This is known as Coulomb’s Law. Lithium is bigger than boron even
though boron has more protons and electrons because the greater number of protons pulls the electrons more
tightly, so the atom shrinks. However, the atomic radius increases (the atoms get bigger) as we go down a
column in the periodic table because we are adding more energy levels. Elements in period 3 are bigger than
elements in period 2 because there are 3 energy levels instead of only 2. So phosphorus is bigger than nitrogen.
Electronegativity is how strongly an atom will attract electrons. Non-metals have higher electronegativities than
metals because non-metals pull on or attract electrons more strongly. Elements in the top right hand corner are
-
more electronegative; fluorine is the most electronegative of all elements. Noble gases are not electronegative
because they do not attract electrons – they have full valence energy levels and do not need any more electrons.
Ionization Energy is the amount of energy required to pull an electron off an atom. Metals lose electrons more
easily because they have a small number of valence electrons, so they have lower ionization energies. Elements
on the bottom of the periodic table have lower ionization energies because the valence electrons are so far away
from the center of the atoms – it is easier to pull an electron off energy level 4 or 5 than it is from 2 or 3.
Sample questions:
1. Describe in your own words the trend (pattern) in atomic radius as you go DOWN a GROUP.
2. Describe in your own words the trend in atomic radius as you go ACROSS a PERIOD.
3. Which of the following elements has the largest atomic radius? Si Ge Sn Pb
a. Why?
4. Describe in your own words the trend in electronegativity as you do DOWN a GROUP.
5. Describe in your own words the trend in electronegativity as you go ACROSS a PERIOD.
6. Which of the following elements has the greatest electronegativity? Cs Ga P O
7. Describe in your own words the trend in ionization energy as you do DOWN a GROUP.
8. Describe in your own words the trend in ionization energy as you go ACROSS a PERIOD.
9. Which of the following elements has the lowest ionization energy? Cs Ga P
O
Quantum mechanics
Quantum Mechanics is the math side of an atom’s energy. A quantum is the amount of energy required to make an
electron jump up one energy level. We can add this energy by heating atoms, like we did in our flame test lab.
When an electron is excited by heat or other energy, it jumps up one energy level. Electrons want to fall back down
to their usual spots, but when they do, they need to release the energy that caused than to jump up one level. The
energy they release is in the form of light in different colors, called the emission spectrum. We can identify
elements based on what color light is released by those electrons falling back to their usual energy levels. For
example, copper shines green light.
Frequency of a wave is how many waves pass a point each second. Wavelength is the distance from the crest of
one wave to the crest of the next wave. As wavelength gets bigger (wider), the frequency decreases. As wavelength
gets smaller (waves are more narrow) the frequency increases.
We can calculate the frequency and wavelength of light by using the formula c=fλ, where c is the speed of light
(3.oox108m/s), f is frequency, and λ (greek letter lambda) is wavelength.
For example, what is the frequency of a wave with wavelength 5.30 x 10-7 m?
c=fλ
3.00 x 108 m/s = f (5.30 x 10-7 m)
divide both sides by (5.30 x 10-7 m)
Be sure to put parantheses around the scientific notation!
3.00 x 108 m/s
f (5.30 x 10−7 m)
=
(5.30 x 10−7 m)
(5.30 x 10−7 m)
f = 5.66 x 1014 Hz
We can calculate the energy of a wave using one of two equations. If we are given frequency, then we can use the
equation E=hf, where E is the energy, h is Planck’s constant (6.63 x 10-34 Js), and f is the frequency. As the frequency
of a wave increases, the energy also increases.
For example, what is the energy of a wave with a frequency of 5.66 x 1014 Hz?
E=hf
E = (6.63 x 10-34 Js) (5.66 x 1014 Hz)
Multiply the two values!
E = 3.75 x 10-19 J
We can also calculate energy of a wave if we are given wavelength. The equation is E=
-34
hc
λ
where E is energy, h is
8
Planck’s constant (6.63 x 10 J-s), c is the speed of light (3.oox10 m/s), and λ (greek letter lambda) is wavelength.
For example, what is the energy of a wave with wavelength 6.7 x 10-7 m?
-34
8
hc (6.63 x 10 Js)(3.oox10 m/s)
E= =
=3.0 ×10-19 J
-7
λ
(6.7 x 10 m)
Sample questions:
1.
2.
3.
4.
5.
What is c? ___________________________
What is h? ___________________________
Do the values of c and h ever change? _________
If wavelength increases, what happens to the frequency? ______________________
If energy increases, what happens to the frequency? ___________________________
6. Red light has a wavelength of 6.9 x 10-7 m. What is its frequency?
7. Blue light has a frequency of 5.66 x 1014 Hz. What is its energy?
8. One type of ultraviolet radiation that tends to result in sunburns has a wavelength of 2.90 x 10-7 m.
Calculate the energy of one photon of that uv radiation.