Unit 10-11 Study Guide
Chemistry
Name _____________________________
1.
Compare Dalton’s, Thomson’s model, Rutherford’s model and Bohr’s model of the atom. (Do each of the
following: (1) Draw a picture of the model, highlighting any new features that were added to this model that the
previous model did not have; (2) What did the individual do learn more about the atom?
Dalton’s Model –
solid sphere
Solid Sphere
Atoms of the same element are identical, Atoms of
different elements are different in properties, The
smallest thing that makes all matter, indivisible, and
atoms can combine to form compounds
Came up with these ideas from studying experiments
from other scientists as well as his own work
Thomson’s Model –
plum pudding or
chocolate chip
cookie model
The atom is made of smaller particles that have a
negative charge (electrons) and since atoms are neutral,
there must also be positive material present.
Rutherford’s Model
-
Did experiments with the cathode ray tube. A beam of
light appears and its path can be altered with a magnetic
field. It must contain mass as light is not affected by a
magnetic field. Only atoms of a gas present inside the
cathode ray tube – conclusion – the negatively charged
particles must come from within the atom
The atom now has a nucleus in the model with electrons
in a space around the nucleus.
Bohr’s Model –
planetary model
Did the Gold Foil Experiment: shot positive alpha
particles at the gold foil and most of the particles passed
through but a few particles would deflect/bounce back.
Concluded that something very small, dense, positively
charged caused the few particles to bounce/deflect and
most of the atom is empty space because most of the
alpha particles passed through the foil.
The atom still has nucleus but now electrons are on
energy levels around the nucleus.
Studied the emission spectrum of the Hydrogen atom.
Each color observed = an amount of energy of light that
corresponds to the energy that an electron emits when
the electron drops from a higher energy level to a lower
energy level.
2.
Which of Dalton’s five principles still apply to the structure of an atom?
*All matter is made of atoms
*Different atoms have different properties
*Atoms can combine to form compounds
3.
Write the symbol for element X, which has 27 electrons and 32 neutrons.
59
27X
(or 5927Co) X-59 (or Cobalt-59)
4.
Determine the number of electrons, protons and neutrons for each of the following:
80
P = 35 E= 35 N= 80-35= 46
35Br
106
46Pd
P = 46 E=46
N= 106-46=60
133
55Cs
P = 55 E=55
N= 133-55=78
5.
How are isotopes of the same element alike?
Isotopes have the same number of protons (and when neutral will have the same number of electrons) so therefore
they are the same element. What is different is the number of neutrons and therefore the mass is different.
6.
Are Ga-70 and Ge-70 isotopes of the same element? Explain your answer.
No – not isotopes as they are different elements (have different number of protons)
7.
Calculate the average atomic mass of magnesium using the following data for three magnesium isotopes.
Isotope
mass (u)
relative abundance
Mg-24
23.985
0.7870
Mg-25
24.986
0.1013
Mg-26
25.983
0.1117
(23.985 × 0.7870) + (24.986 × 0.1013) + (25.983 × 0.1117) = Answer
18.876 +
2.531 +
2.902
=
Answer
24.309 = Answer
8.
Answer the following questions:
a. How many orbitals are in a d sublevel? 5
b. What is the maximum number of electrons that can be contained in an f sublevel? 14
c. How many sublevels are in the 2nd energy level? 2
d. How many energy levels are represented on the periodic table? 7
9.
Describe/draw the general shape of the orbitals in the s, p, and d sublevel. How many of each kind of
orbital are present in the s, p, d and f sublevels?
s
p
d
f
1
10.
3
5
7
Define: Electron configuration – representation of the arrangement of electrons that are distributed among
the energy levels, sublevels and orbital shells.
Orbital notation – A visual transformation of the electron configuration. Shows where each
specific electron is placed (in order) and it’s spin.
11.
Answer the following questions:
a.
Which rule/principal explains that there can only be two electrons in an orbital and they must be in
opposite spins? Pauli Exclusion Principle
b.
Which rule/principal explains that there is an order to how the electrons can fill energy levels?
Aufbau principle
c.
Which rule/principal explains that when similar energy sublevels are filled, each orbital must have
one electron before receiving two? Hund’s Rule
12.
Use the Aufbau principle, Pauli exclusion principle and Hund’s rule to write electron configurations and
orbital notations. (Know how to do Noble Gas and extended electron configs)
a.
Fluorine 1s2 2s2 2p5
[He] 2s2 2p5
1s2 2s2 2p6 3s2 3p6 4s2 3d10
[Ar] 4s2 3d10
b.
Zinc
c.
Sulfur 1s2 2s2 2p6 3s2 3p4
[Ne] 3s2 3p4
d.
Tin
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p2
[Kr] 5s2 4d10 5p2
13.
What are the letters that are assigned to each of the four quantum numbers? What does each of the quantum
numbers describe about an electron?
Principle Quantum
Angular Momentum
Magnetic Quantum
Spin Quantum Number
Number
Quantum Number
Number
n
l
m
s
Energy level the
Sublevel the electrons
Which orbital in the
Spin – indicating that no
electrons are on
are in (s, p, d or f)
sublevel (which
two electrons can be in
arrangement in space of
the same energy level,
the # of shapes for the
sublevel and orbital and
sublevel)
have the same spin
14.
As an electron in an excited state returns to its ground state,
a.
light energy is emitted
b.
energy is absorbed by the electron
c.
the atom is likely to undergo spontaneous decay
d.
the electron configuration of the atom changes
15.
According to quantum theory, an electron
a.
remains in a fixed position
b.
can replace a proton in the nucleus
c.
occupies the space around the nucleus only in certain well-defined orbitals
d.
has a certain velocity and position, of which both are known
16.
We did a lab called the Flame Test Lab…
a.
What part of the atom caused the various colors we saw?
electron
b.
Why were there different colors for different elements?
Different number of electrons on a variety of energy levels
c.
Knowing that these sticks were soaked in water solutions of various metal compounds, what would
be the difficulty in identifying metals dissolved in drinking water? Too many different metals to
distinctly see any one color or too low concentration of metals to see any color at all.
17.
Why are Hydrogen and Helium set off from the rest of the periodic table? Discuss each one individually.
Hydrogen has only one valence electron like other elements in group 1A. Hydrogen can lose that one electron to
form a +1 charged ion like others in the Alkali metal group BUT…it is not a metal and doesn’t have any properties of
a metal. Helium is unreactive and stable like elements in group 8A/18. Its outermost energy level is full BUT…it
only has two valence electrons, not eight.
18.
How does electron configurations relate to its position on the periodic table? (i.e. group number and period
number)
The electron configuration tells you where an element is located on the P.T. The period number is the highest
principle Quantum number and the group # (#A) tells how many valence electrons there are. Example:
1s2 2s2 2p6 3s2 3p2 – The highest principle Q# is 3 – the valence electrons are in the 3rd energy level, the element is in
the 3rd period. The highlighted superscripts = 4. There are four valence electrons – it is in group 4A.
19.
What do elements in the same group have in common? What do elements in the same period have in
common?
Elements in the same group have the same ending electron configuration (same number of valence electrons) and for
that reason they all behave in a similar fashion/have similar physical characteristics. Elements in the same period
have their valence electrons in the same outermost energy level.
20.
If an element ends with an electron configuration of #s2 #p4…
a. What group number is this element found? Group 6A/16
b. How many valence electrons does it have? 6 valence electrons
c. Would it gain or lose electrons to be stable? Gain electrons to get 8 valence electrons
d. How many electrons will it lose or gain? Gains two electrons
e. What will the ion charge be (number and +/- value)? -2 Charge
f. If this element has its valence electrons in the 3rd energy level, what element is it? Sulfur
21.
If an element has an ending electron configuration of 4s23d104p5…
a. What period is the element in? 4th period
b. How many valence electrons does it have? 7 valence electrons
c. Which element has this electron configuration? Bromine
22.
What do atoms that have similar properties share in common regarding their electron structure?
Elements with similar properties have the same number of valence electrons / ending e-configuration.
23.
Which is larger; the Chlorine atom or the Chlorine ion? WHY?
The Chlorine ion – the ion has more electrons than the neutral atom for the same charged nucleus.
24.
Why are cations smaller than their neutral counterparts?
Cations have fewer electrons (lesser # of energy levels of electrons) for the same number of protons
Why does the atomic radius increase as you go down a group?
Radius of an atom increases due to more energy levels of electrons being added to the next atom in the group.
25.
Why does atomic radius decrease as you go across a period?
Radius decreases going across a period because although there are more electrons being added, they are added to the
same energy level (not more energy levels) but the nucleus is also having more protons (stronger nuclear charge) to
pull in the electrons from the same energy level.
26.
Why does ionization energy increase as you go across a period?
Ionization energy increases going across a period because as the element gets smaller and as the nucleus holds on to
the valence electrons tightly, more energy is needed to take away electrons. In addition, the elements on the right are
approaching stability or are stable and do not want to lose any electrons.
27.
In each combination listed, indicate which element has the higher electro negativity and why?
Sulfur + Aluminum Sulfur (it is further right, smaller atom, tighter hold on electrons)
Oxygen + Selenium Oxygen (it is closer to the top/a smaller atom, tighter hold on electrons)
Gallium + Phosphorus Phosphorus (it is further right, smaller atom, tighter hold on electrons)
Bromine + Chlorine Chlorine (it is closer to the top/a smaller atom, tighter hold on electrons)
28.
Circle the element that meets the requirement of each question that is asked:
a) Which Metal is MORE Reactive?
Na
or
Al
d) Which is GREATER in Size (dealing with IONS)?
K+1
or
K
K
or
Ca
S2-
or
S
K
or
Rb
Mg+2
or
Al+3
Mg
or
Ba
Ba+2
or
Ca+2
K
or
Mg
Cl1-
or
I1-
N3-
or
Br1-
or
Cl1-
b) Which has a LARGER Radius?
Mg
or
S
S2-
Br
or
Cl
e) Which has the GREATER Electronegativity?
Al
or
Ca
O
or
C
Al
or
P
c) Which has a LARGER Ionization Energy?
O
or
Ne
S
or
Se
Li
or
N
S
or
Mg
P
or
As
Na
or
N
K
or
S
e) Which Nonmetal is MORE reactive?
F
or
Br
Cl
or
I
P
or
Cl
29.
What is the reason for the trend in electronegativity as one goes across the period? Why do noble gases have
no values for electronegativity?
Electronegativity is the attraction an element has to shared electrons when bonded to another atom. Electronegativity
increases going across a period. This occurs because atoms near the right side are smaller atoms; the nuclear charge
is stronger thus pulls in the shared electrons more. Noble gases have no value for elecronegativity because they do
not form bonds with other atoms; they are stable and nonreactive.
30.
What is the shielding effect and how does it affect electronegativty and ionization energy?
Shielding effect is the decrease in attraction between valence electrons and the nucleus due to an atom’s inner
electrons. Since electronegativity is the attraction for electrons and ionization energy is the energy needed to remove
an electron for form an ion, these both relate to shielding effect because as an atom’s size increases (more inner
layers of electrons), the greater the shielding effect, the lower the electronegativity and the lower the energy needed
to take away electrons.