Download Metal found in the salt

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Unit 2 Bonding & Atomic Structure
Syllabus:
Lesson
2.1
Homework
Pg 30
2.2
Description
Basic Atomic Structure Guided Notes: Subatomic Particles,
Rutherford Experiment, Atomic Target Practice Activity
Ions, Isotopes, and Weighted Atomic Mass
2.3
General Properties of the Periodic Table
Pg 31
2.4 & 2.5
Electron configuration & Quantum Numbers
Pg 32-33
2.6
Energy Calculations and Flame Test Lab
Pg 33
2.7
Period Trends
Pg 34
2.8
Periodic Trends Activity
Pg 34
2.9
Exam 2
Pg 35
Pg 31
Atomic Theory
 Democritus
 John Dalton: identify Dalton’s postulates and evaluate their accuracy in light of current atomic
theory.
 JJ Thomson: explain how Thomson’s experiments led to his discovery of electron properties.
 Ernest Rutherford: explain how Rutherford’s experiments led to his conclusions about the
nuclear atom.
 Niels Bohr: explain Bohr’s addition of electron energy levels to Rutherford’s nuclear atom.
 Erwin Schrodinger

James Chadwick
Properties of Periodic Table
 explain how the contributions of scientists such as Mendeleev led to the development of the
Periodic Table as a predictive tool
 explain how chemical and physical properties were historically used to develop the Periodic
Table
 identify and explain properties of chemical families, including alkali metals, alkaline earth
metals, halogens, noble gases, and transition metals
 use the Periodic Table to identify the chemical family of an element and explain its properties
Periodic Trends: identify and explain periodic trends, including:
 atomic radius
 ionic radius
 electronegativity
 ionization energy
 electron affinity
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Electron Configuration
 express the arrangement of all electrons in an atom using electron configuration
 express the electron configuration of the d-block element
 using the idea of quantum mechanics describe the areas of probability where an electron can be
located and assign the four quantum numbers
 from the electron configuration identify which electrons are valence electrons and how they are
used in bonding
Electromagnetic Spectrum
 describe the electromagnetic spectrum
 describe the mathematical relationships between energy, frequency, and wavelength of light
 calculate wavelength and frequency using the speed of light
 calculate energy and frequency using Planck’s constant
Unit 2 Vocabulary
 quantum
 quantum mechanical
model
 electron
configuration
 energy level
 sublevel
 atomic orbital
 aufbau principle
 Hund’s rule
 Pauli exclusion
principle
 Shell
 Subshell
 Block



















Orbital notation
Electron
configuration
Bohr model
periodic law
chemical
family/group
representative
element
alkali metal
alkaline earth metal
transition metal
inner transition
metal
halogen






Atomic radius
Ionic radius
Electronegativity
Ionization energy
Group trend
Periodic trend
Valence electrons
electromagnetic
radiation
spectrum
wavelength (𝛾)
frequency (𝜐)
hertz (Hz)
Planck’s constant (h)
speed of light (c)
2.1 Basic Atomic Structure Guided Notes
What is an Atom???
 Atoms are NOT indivisible! They CAN be broken down into smaller particles.
Subatomic Particle
Proton
Neutron
Electron
Charge
Mass
Location
Protons, neutrons and electrons are __________evenly distributed in an atom.



1 amu (atomic mass unit)= 1.66 x 10-27 kg (also: 1/12th the mass of a carbon atom)
The protons and neutrons exist in a _________________ core at the center of
the atom. This is called the ___________________. The nucleus has a
_______________________ charge because the protons have a positive charge
and the ______________________ don’t have a charge!
The _____________________________ are spread out around the edge of the
atom. They orbit the nucleus in layers called ____________________. Electrons
have a _________________________ charge.
Drawing of an
Atom
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
The ___________________ of an element contain equal numbers of protons and electrons and so have
no overall charge, so if you can find it on the Periodic Table, it means it has a charge of
__________________!!!
Atomic Number (Z)
 The atoms of any particular element always contain the same number of _________________. For
example: hydrogen atoms always contain ____ proton and carbon atoms always contain _____ protons.
 The number of protons in an atom is known as the ______________ _______________.
 It is the ______________________ of the two numbers shown in most periodic tables.
 If the number of protons changes, then the atom becomes a different ____________________.
 Changes in the number of particles in the nucleus (protons or neutrons) are very _________. They only
take place in ___________________ processes such as: radioactive decay, nuclear bombs or nuclear
reactors.
Mass Number (A)
 Electrons have a mass of almost _________________, which means that the mass of each atom results
almost entirely from the number of ______________ and neutrons in the nucleus.
 The ____________ of the protons and neutrons in an atom’s nucleus is the __________________
number. It is the _________________________of the two
numbers shown in most periodic tables.
Mass Number= # Protons + # Neutrons
 That means in order to figure out the number of neutrons, you
simply use the following:
# Neutrons =Mass # - # Protons (Atomic #)
Electrons

Atoms have no overall electrical charge and are ____________________.
 This means atoms must have an equal number of _________________ and ____________

The number of electrons is therefore the same as the ___________________ number.

Atomic number is the number of ____________________rather than the number of electrons,
because atoms can lose or gain electrons but do not normally lose or gain protons.
2.1b Atomic Target Activity: Rutherford Scattering
Introduction:
The Rutherford gold foil experiment is one of the most famous of all time. More than 25 years
after conducting the experiment, Ernest Rutherford described the results this way:
“It was about as credible as if you had fired a 15-inch shell
at a piece of tissue paper and it came back and hit you.”
The experiment itself was actually the culmination of a series of experiments, carried out over a five-year
period, dealing with the scattering of high-energy alpha particles by various substances.
Ernest Rutherford received the Nobel Prize in Chemistry in 1908 for his investigation into the
disintegration of the elements as a result of radioactive decay. Among the products of the radioactive
decay of elements are alpha particles—small, positively charged, high-energy particles. In trying to learn
more about the nature of alpha particles, Rutherford and his co-workers, Hans Geiger and Ernest
Marsden, began studying what happened when a narrow beam of alpha particles was directed at a thin
piece of metal foil. Alpha particles are a type of nuclear radiation, traveling at about 1/10 the speed of
light. As expect4ed for such high-energy particles, most of the particles penetrated the thin metal foil and
were detected on the other side. What was unexpected was that a very few of the alpha particles were
actually reflected back toward the source, having been “scattered” or bent due to their encounters with
the metal atoms in the foil target. The number of alpha particles that were reflected back depended on
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the atomic mass of the metal. Gold atoms, having the highest atomic mass of the metals studied, gave
the largest amount of so-called “backscattering”
Rutherford’s scattering experiments have been described as a “black box” experiment. The
properties of the alpha particles, their mass, charge, speed, etc., were at least partially understood. The
atoms making up the target, however, presented Rutherford with a kind of black box; the structure of the
atom was not known at the time. In order to explain the results of the scattering experiment, Rutherford
had to propose a new model of the atoms. A model that explained the results of the data gathered from
the experiment. In 1911 Rutherford proposed the following model for the structure of the atom:



Most of the mass of the atom is concentrated in a very small, dense central area, later called the
nucleus, which is about 1/100,000 the diameter of the atom.
This was proposed as a result of what data:
_____________________________________________________________________________
_________________________________________
The rest of the atom is apparently “empty space”
This was proposed as a result of what data
_____________________________________________________________________________
_________________________________________
The central, dense core of the atom is positively charged, with the nuclear charge equal to about
one-half the atomic mass.
This was proposed as a result of what data
_____________________________________________________________________________
_________________________________________
Objective:
The purpose of this activity is to discover by indirect means the size and shape of an unknown
object, which is hidden underneath the middle of a large board. By tracing the path the marble takes after
striking the unknowntarget from a variety of angles, it should be possible to estimate the general size and
shape of the unknown target.
Pre-Lab Questions
1. Read the material in your textbook about the Rutherford experiment and answer the three
questions in the introduction.
2. This activity is a simulation of Rutherford’s scattering experiment. Read the entire procedure and
compare the components used in this simulation to Rutherford’s original discuss what each
component in our simulation corresponds to in the original experiment.
3. The key skills in this activity, as in Rutherford’s experiment, are the ability to make careful
observation and to draw reasonable hypotheses. Assume that the marble strikes following sides
of a possible target. Sketch the path the marble might be expected to take in each case.
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4. Discuss what information can be inferred if the marble rolls straight through without striking the
unknown target.
Materials:
1. foam board with unknown shape attached
2. marbles
3. white paper
4. push pins
5. pencil
6. ruler
Procedure:
1. Form a group of three students
2. Pin the paper to the top of the board (do not look at the shape on the underneath side)
3. Roll the marble with a moderate amount of force under one side of the board. Observe where the
marble comes out and trace the approximate path of the marble on the paper.
4. Working from all four sides of the board, continue to roll the marble under the board, making
observations and tracing the rebound path for each marble roll. Roll the marble AT LEAST 20
TIMES from each side of the box. Be sure to vary the angles at which the marble is rolled. You
may use the rulers as a launching platform.
5. After sketching the apparent path from all sides and angles, the general size and shape of the
unknown target should emerge.
6. Form a working hypothesis concerning the structure of the unknown target. Based on this
hypothesis, repeat as many “targeted’ marble rolls as necessary to confirm or revise the
structure.
7. Check your answer with your teacher. DO NOT look under the board.
8. If time permits try an extension, or another shape.
Post Lab Questions
1. Draw the general size and shape of the target to approximate scale in the square below.
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2. The speed of the marble rolls was an uncontrolled variable in this activity. How would the
outcome of the scattering test have been different if the marble speed had been faster or slower?
3. Compare the overall size of the target with the size of the marble used to probe its structure.
How would the outcome of the scattering test have been different if different size marbles had
been used? Explain.
2.2 Ions, Isotopes, and Weighted Atomic Mass
Isotopes
A) Atoms with same number of protons ( #p+ ) but different numbers of neutrons (#no). Isotopes have the same
atomic number (Z) but different mass numbers.
B) Recognizing Isotope Notations
i) X = element from periodic table
ii) A = mass number = #p + #n
iii) Z = atomic number = #p
-NOTE: In a neutral atom, Z = #p = #e
C) Examples
Sodium-23 Na-23
(Z=11)
Sodium 24
(Z= 11) 11 p, 13 n, 11 e
Na-24
11 p, 12 n, 11 e
Ions
A) Ions are charged atoms.
B) Charges arise due to loss or gain of electrons.
C) If atom loses electrons, a positively charged cation is formed.
Example: Li ──> Li+ + 1 eD) If atom gains electrons, a negatively charged anion is formed.
Example: F + 1 e- ──> FE) In ordinary matter, cations and anions always occur together so that matter is charge-neutral overall.
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Atomic Mass
1) Atomic Mass of Element represents the weighted average of all the naturally occurring isotopes of that
element.
2) Value found below the atomic symbol of the element in the periodic table.
A) Calculating Atomic Mass for an Element
i)
ii)
iii)
iv)
Must be given relative abundance (%) for each isotope
Must know isotopic mass for each isotope
Convert percent abundance to a decimal (divide by 100%)
Use following formula to obtain result
Atomic Mass = Σ (fraction of isotope n) x (mass isotope n)
Example:
Find Atomic Mass of Carbon given the following data.
Data:
98.89% Carbon 12
1.11% Carbon 13
isotopic mass 12 amu
isotopic mass 13.0034 amu
Solution: (12 amu * 0.9889) + (13.0034 amu * 0.0111)
=
11.867 amu +
0.144 amu = 12.011 amu
Atomic Weights/Mass
Naturally occurring magnesium consists of three stable isotopes:
isotope
amu
Abundance
Mg-24
23.985
78.99%
Mg-25
24.986
10.00%
Mg-26
25.983
11.01%
What is the atomic weight of Magnesium?
Ans:_______________________
Naturally occurring silicon consists of three stable isotopes:
isotope
amu
Abundance
Atomic weight = 28.09 amu
Si-28
27.977
92.21%
Si-29
28.976
4.70%
Si-30
?
3.09%
What is the atomic mass of Si-30?
Ans:_______________________
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Protons, Neutrons, Electrons
Fill in the table below with the correct numbers (first one is done as an example)
Symbol
name
Atomic
Mass
charge
# of parts
# of
number
number
in nucleus
protons
23
11
Na
39
19
K
41
19
K
41
19
K 1
12
25
# of
neutrons
# of
electrons
0
1-
35
18
7
7
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2.3 General Properties of the Periodic Table
Mendeleev's Periodic Table (1869)
A. Organization
1. Vertical columns in atomic weight order
a. Mendeleev placed elements in rows with similar properties
2. Horizontal rows have similar chemical properties
B. Missing Elements
B . Gaps existed in Mendeleev’s table
a. Mendeleev predicted the properties of the “yet to be discovered” elements
(1) Scandium, germanium and gallium agreed with his predictions
C. Unanswered Questions
1. Why didn't some elements fit in order of increasing atomic mass?
2. Why did elements exhibit periodic behavior?
Moseley and the Modern Periodic Table (1911)
A. Protons and Atomic Number
1. The periodic table was found to be in atomic number order, not atomic mass order
B. The Periodic Law
1. The physical and chemical properties of the elements are periodic functions of their atomic
numbers
2. Elements with similar properties are found at regular intervals within the periodic table
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* Moseley was killed in battle in 1915, during WWI. He was only 28 years old
Organization of the Table
1. Groups or Families
a. Vertical columns containing elements with similar chemical properties
2. Periods (series)
a. Horizontal rows of elements
3. Metals and Nonmetals
a. A stair-step line on the table separates the metals from the nonmetals
b. Metalloids (Semimetals) straddle the line and have properties of both metals and
nonmetals
4. Lanthanide and Actinide Series (Inner Transition Metals)
a. Metals and man-made metal elements
5. Group 1 – Alkali metals (the most reactive metal elements) (except hydrogen (H) also in this
group)
6. Group 2 – Alkaline earth metals (very reactive metal elements)
7. Group 17 – Halogens (the most reactive nonmetal elements)
8. Group 18 – Noble gases (the least reactive elements – inert and very stable)
Types of Elements
A. Metals
1. Luster
2. Good conductors of heat and electricity
3. Malleable
4. Ductile
5. High tensile strength
B. Nonmetals
1. Many nonmetals are gases at room temperature
2. Solid nonmetals tend to be brittle and non-lustrous
3. Poor conductors of heat and electricity
C. Metalloids
1. Some properties of metals and some properties of nonmetals
2. Solids at room temperature
3. Semiconductors of electricity
D. Noble Gases
1. All are gaseous members of group 18
2. Generally unreactive and stable
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Periodic Table
• Prior to 1860 no agreement/method to
accurately determine masses of atoms.
• First International Congress of Chemists –
1860
– Stanislao Cannizzaro presented method for
accurately measuring atomic masses
– Looked for relationships between atomic
masses and other properties of elements
Julius Lothar Meyer
John Newlands
• Noticed elements properties repeated
every 8th element when arranged by
atomic mass
• Named this phenomenon “the Law of
Octaves”
• Did not work for all elements
• First tables arranged elements by atomic
weight
– Could not agree on atomic weights therefore
tables were different
• Developed first modern
table
– Consisted of 28 elements
divided into 6 families
– Families (groups) had similar
chemical and physical
properties
– Discovered all elements in
same family had same
number of valence e- -outermost electrons in
highest energy level
– Why?
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Dmitri Mendeleev
• Noticed that properties
repeat themselves at
certain intervals
• Arranged all known
elements into one table
based on properties– 1869
• 1871 - Proposed the
“Periodic Law”
• Based on the properties
spaces were left for
unknown elements (Sc,
Ga, Ge)
• Upon discovery of other elements
inconsistencies were found with
Mendeleev’s table
• Atomic masses improved and they no
longer arranged the elements by
increasing atomic mass
• Why can most elements be arranged by
atomic mass?
• What was the reason for chemical
periodicity?
Henry Mosely
• Discovered elements contain unique
number of protons (atomic number) - 1911
• Arranged elements by atomic number 1913
• Fully explained the Periodic Law
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2.4 Quantum Numbers & Electron Configuration
Principal Quantum Number (n)
• Indicates the relative sizes and energies of
atomic orbitals
• n is an integer greater than 1 (n =1, 2, 3,….)
• 7 energy levels have been identified for
Hydrogen
Quantum Numbers and Electron
Configuration
1A
2A
3B
4B
5B
6B
7B
8B
8B
8B
1B
2B
3A
4A
5A
6A
7A
8A
1A
Quantum Mechanical Model
• Electrons are attracted to the nucleus by
electrostatic forces between oppositely
charged objects
• Electrons reside in space that are different
distances from nucleus
• Limited number of regions where an e- can
reside (energy is quantized)
• Atoms absorb/emit radiation when e- move
Continued…
• Energy states have negative values
• Energy values increase (become more
positive) farther from nucleus n= ∞ has 0
energy value
• Can completely remove an electron from an
atom when n=∞ (an ion is formed)
2A
3B
4B
5B
6B
7B
8B
8B
8B
1B
2B
3A
4A
5A
6A
7A
8A
Angular Momentum Quantum # (l)
• Describes the shape of the
orbitals in the sub level
• Principal levels (shells)
contain sublevels
• Number of sublevels equal to
n
• Number of orbitals in
sublevel is always an odd
number
• Sublevels are designated s, p,
d, & f depending on the
shape of the sublevel
l
Letter
0
s
1
p
2
d
3
f
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1A
2A
3B
4B
5B
6B
7B
8B
8B
8B
1B
2B
3A
4A
5A
6A
7A
d - subshell
8A
s
• d – subshell contains 5 orbitals
p
d
f
f - subshell
Magnetic Quantum Number (ml)
• Describes the orientation of an orbital around
the nucleus
• s orbital is spherical and has only 1 orientation
(m = 0)
• p orbital can orient along each axis (x, y, &
z)(m = -1, 0, +1)
• f – subshell contains 7 orbitals
s - subshell
• s – subshell contains one orbital
1A
p - subshell
• p – subshell contains 3 orbitals
2A
3B
4B
5B
6B
7B
8B
8B
8B
1B
2B
3A
4A
5A
6A
7A
8A
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Pauli Exclusion Principle
Spin Quantum # (ms)
• spin makes the electron behave like a tiny
magnet
• spin can be clockwise or counterclockwise
• spin quantum number can have values of +1/2
or -1/2
• No two electrons in the same atom can have
the same set of four quantum numbers.
• Electrons must have opposite spins in the
same orbital
• Spins of electrons represented with
• Can only put 2 electrons in each orbital and
must have opposite spins
Hund’s Rule
Predicting Electron Configuration
• Atoms like to have the most stable configuration as
possible.
• Number of subshells equal to shell number, n
• In order of increasing energy subshells labeled s, p, d,
&f
4s < 4p < 4d < 4f
• Always odd number of orbitals in a subshell
• Maximum number of e- in subshell equal 2x number
of orbitals
• Electrons are added to an atom, one at a time,
starting with lowest available orbital – Aufbau
principle
1A
2A
3B
4B
5B
6B
7B
8B
8B
8B
1B
2B
3A
1s
2
p
2s
3s
3p
4s
3
d
4p
5s
4d
5p
6s
5d
6p
7s
6d
7p
4
f
4A
5A
6A
7A
• Orbitals of equal energy are each occupied by
one electron before any orbital is occupied by
a second electron.
• All electrons in a singly occupied orbital must
have same spin
Orbital Notation
8A
• Unoccupied orbital is designated with a
_____ and the orbitals name written
underneath
3p
3p
3p
• Electrons are placed in each orbital using
• Give the orbital notation for carbon and
fluorine
5f
Carbon
• 6 electrons
1s
2s
2p
15
fluorine
• 9 electrons
Short Hand Notation
• Locate the Noble gas preceding your element
– The noble gas has a full outer shell in its electron
configuration
• Place the symbol for the Noble gas in brackets
– [X]
1s
2s
2p
• Complete the electron notation for the
desired element
Electron Configuration Notation
• Number of electrons in a sublevel is shown by
adding a superscript to the sublevel
designation
• Boron has 5 electrons
– 1s22s22p1
• Give the proper configurations for carbon and
fluorine
Example
• Calcium
1s
•
•
•
•
2s
2p
3s
3p
4s
Noble Gas preceding Ca is Argon (Ar)
[Ar]4s2
Give short hand notation for Bromine
[Ar]4s23d104p5
Summary video on electron
configuration
• Carbon: 1s22s22p2
• Fluorine: 1s22s22p5
• http://www.youtube.com/watch?v=Vb6kAxwS
WgU
Electron Configuration Practice Worksheet
In the space below, write the unabbreviated electron configurations for the following elements:
1)
Sodium
________________________________________________
2)
Iron
________________________________________________
3)
Bromine
________________________________________________
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In the space below, write the abbreviated electron configurations for the following elements:
6)
Cobalt
________________________________________________
7)
Silver
________________________________________________
8)
Tellurium
________________________________________________
Determine what elements are denoted by the following electron configurations:
11)
1s22s22p63s23p4 ____________________
12)
1s22s22p63s23p64s23d104p65s1 ____________________
13)
[Kr] 5s24d105p3 ____________________
Determine which of the following electron configurations are not valid:
14)
1s22s22p63s23p64s24d104p5 ____________________
15)
1s22s22p63s33d5 ____________________
16)
[Ra] 7s25f8 ____________________
2.5 Energy Calculations & Flame Test Lab
Wave Description of Light
• Electromagnetic radiation (ER): form of
energy that exhibits wavelike behavior as it
travels through space
• All forms of ER move at a constant speed of
about 3.0 x108 m/s (speed of light, c)
• Wavelength (λ): distance between
corresponding points on adjacent waves.
• All forms of ER together make the
electromagnetic spectrum
• Frequency (ν): number of waves that pass a
specific point in a given time, usually one
second.
• Unit: Hertz (Hz), aka (1/s) or (s-1)
• For electromagnetic radiation, frequency and
wavelength are related
C=λν
• If λ increases, what must happen to ν? Does c
change?
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Particle description of light
• Max Planck, 1900s, suggested that object emit
energy in small, specific amounts called
quanta
• Quantum: minimum amount of energy that
can be gained or lost by an atom
• 1905 Einstein expands on this idea. ER have
dual wave/particle nature.
• While light emits many wavelike particles, it
can also be thought of as a stream of those
particles
• Einstein named the particles photons
• Photon: particle of ER having zero mass and
carrying a quantum of energy
The Hydrogen Atom
• Ground state: lowest energy state of an
atom
• Excited state: atom has higher potential
energy than ground state
• When an excited atom returns to ground
state, it gives off energy it gained in the
form of ER.
• Production of colored signs (neon) is
example
• Planck proposed a relationship between a
quantum of energy and the frequency of
radiation
E=hν
• E is energy, in Joules, of a quantum of
radiation
• h is Planck’s constant (fundamental physical
constant)= 6.626 x 10-34 J*s
• ν is freqency of radiation
• Energy of a particular photon depends on the
frequency of the radiation
Ephoton= hν
• Einstein’s Explanation: ER is absorbed by matter
only in whole numbers of photons
– For e- to be ejected, must be struck by single photon
possessing at least minumum energy
– According to equation, this energy corresponds to
frequency
– If photon’s frequency is below minimum, no e- ejected
Experiment
• Pass electrical current through H gas in
vacuum tube at low pressure
• Emits characteristic pink glow
• When the light was passed through prism,
separated into series of specific frequencies
and therefore wavelength (… equation?) of
visible light
• These bands are Hydrogen’s line emission
spectrum
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Implications
• Since H atoms emit only specific frequencies,
difference between energy states must be
fixed.
• Therefore, e- of H atom exists only in very
specific energy states
• In 1913, Bohr proposed a model that linked
the atom’s electron with photon emission
• Energy is higher in orbits farther from
nucleus (like a ladder)
Problem?
• Classical theory predicted that H atoms would
be excited by whatever amount of energy was
added to them.
• Expected to observe continuous range of
frequencies of ER, or continuous spectrum.
• Why had H only given off specific frequencies
of light?
•  Quantum Theory
• Excited H atom falls back from excited state to
ground state, and emits photon
• This energy is equal to difference between
initial and final state
• Based on the wavelengths of hydrogen’s line
emission spectrum, Bohr calculated energies
the e- would have in the allowed energy level
for H atom.
• Bohr’s calculated values agreed with
experimentally observed values for lines in
each series
• Scientists tried to apply this model to other
element’s atoms
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Energy Calculations Practice:
1. What is the energy of a mole of photons with a wavelength of 1.60 10-3 m?
2. What is the energy of a photon with a wavelength of 2.65 10-4 m?
3. What is the frequency of a photon with a wavelength of 3.70 10-6 m?
4. What is the frequency of a photon with a wavelength of 8.60 103 nm?
2.6b FLAME TEST LAB
In this activity, you will investigate the colors of flame produced by solutions of metal salts. When a substance is
heated in a flame, the atoms absorb energy from the flame. This absorbed energy allows the electrons to be
promoted to excited energy levels. From these excited energy levels, there is a natural tendency for the electrons
to make a transition or drop back down to the ground state. When an electron makes a transition from a higher
energy level to a lower energy level, a particle of light called a photon is emitted. Both the absorption and
emission of energy are quantized – only exact amounts of energy are required.
An electron may drop all the way back down to the ground state in a single step, emitting a photon in the process.
Alternatively, an electron may drop back down to the ground state in a series of smaller steps, emitting a photon
with each step. In either case, the energy of each emitted photon is equal to the difference in energy between the
excited state and the state to which the electron relaxes. The energy of the emitted photon determines the color
of light observed in the flame.
In this activity, metal salts in alcohol are burned, producing different colored flames. By comparing the color given
off by an unknown with the known metal salts, the identity of the metal salt can be determined.
Flame Tests Activity
Materials:
 matches
 Spray bottle(s) from your teacher
 ethanol
 the following metal salts
 lithium chloride
 strontium chloride





calcium chloride
copper(II)chloride
sodium chloride
potassium chloride
UNKNOW
20
Procedure:
1. Obtain a spray bottle of each metal ion.
2. Light your Bunsen burner.
3. If your flame blows out TURN OFF the gas immediately and relight the burner.
4. Holding the bottle upright and approximately 12 inches from the flame squirt the contents at the flame. It
may be necessary to move the location and the distance of the squirting depending on air currents in the
room.
5. Record the color of the resulting flame and the intensity, be descriptive in your colors.
6. Repeat for all the salts.
Data Table:
Metal found in the salt
Lithium
Strontium
Calcium
Copper
Sodium
Potassium
UNKNOWN
Identity:_____________
Flame Color and Intensity
21
Flame Test Analysis
1. List the colors observed in this lab from the highest energy to the lowest energy.
2. List the colors observed in this lab from the highest frequency to the lowest frequency.
3. List the colors observed in this lab from the shortest wavelength to the longest wavelength.
4. What is the relationship between energy, frequency, and wavelength?
5. Do you think we can use the flame test to determine the identity of unknowns in a mixture?
Why or why not?
6. How are electrons “excited” in this part of the experiment? What does it mean when the
electrons are “excited”?
7. Explain why we did not see distinct lines (like on an emission spectrum) when the metal salts
were burned.
8. Why do different chemicals emit different colors of light?
9. Why do you think the chemicals have to be heated in the flame first before the colored light is
emitted?
10. Colorful light emissions are applicable to everyday life. Where else have you observed colorful
light emissions? Are these light emission applications related? Explain.
22
2.7 General Period Trends
Trends in the
Periodic Table
s-Block Elements
• Groups 1 & 2
• All elements in group 1 & 2 will have an
electron configuration of
– ns1 or ns2 where n = highest energy level
occupied
Periodic Law
• The physical and chemical properties of
the elements are periodic functions of their
atomic numbers.
• Aka – when elements are arranged by
increasing atomic number, elements with
similar properties appear at regular
intervals.
Parts….
•
•
•
•
•
Alkali metals – group 1
Alkaline earth metals – group 2
Halogens – group 17
Transition metals – d block elements
Inner Transition metals
– Lanthanides (elements 58-71) added in early
1900’s
• Have very similar properties
– Actinides (elements 90-103)
Alkali Metals
• Group 1 elements
• In the elemental state
– Soft
– Silvery metal
– High melting points
– Extremely reactive therefore are not found in
elemental state in nature
• React violently with water to produce
hydrogen gas
Alkaline – Earth Metals
• Group 2 elements
• Outer most s orbital is full
– Do not exhibit stability (outer p orbital is
empty)
• Properties
– Harder, denser than group 1
– Higher melting points than group 1
– Not as reactive but too reactive to be found in
nature in elemental form
23
Burning Mg
p-block elements
• All elements in p block have a full s orbital
• Properties
– Contain all non metals except H & He
– Contain all metalloids (exhibit properties of both
metals and non metals)
• Have semi conducting properties
– Contains 6 metals
• Elements in s & p block make up the
representative elements
Halogens
Hydrogen & Helium
• H has same valence electrons as group 1
but does not share any other properties
• He share same electron configuration
(valence e-) as group 2 but does not share
same properties
– Placed with group 18 because it is very stable
• Group 7A/17
– Most reactive non metals (Fluorine is most
reactive)
– Will bond with a metal to form a salt
– F & Cl are gases at room temp
– Br is a liquid at room temp
– I & At are solids at room temp
Trends in Atomic Size
Exceptions in the d-block
• The following elements have odd
configurations
– Cr: [Ar]4s13d5
– Cu: [Ar]4s13d10
– Ag: [Kr]5s14d10
• More stable with half filled s & d orbitals or
full d orbital
• Exceptions follow throughout the d
element similar to Chromium and Copper
d-block elements
• Transition elements
– Beginning filling the 3d orbitals
– Good conductors of electricity
– High luster
– Less reactive than s-block elements
• Can be found in elemental form
• - Period - atomic radius decreases as you go
from left to right across a period.
• Why? Stronger attractive forces in atoms (as you
go from left to right) between the opposite
charges in the nucleus and electron cloud cause
the atom to be 'sucked' together a little tighter.
Atomic Radius
• Atomic radius is simply the radius of the atom,
an indication of the atom's volume.
• Atomic radius is one-half the distance between
the two nuclei in a molecule consisting of two
identical atoms.
24
Trends in Atomic Size cont.
• Group - atomic radius increases as you go down a
group.
Why?
• There is a significant jump in the size of the nucleus
(protons + neutrons) each time you move from period to
period down a group.
• Additionally, new energy levels of elections clouds are
added to the atom as you move from period to period
down a group, making the each atom significantly more
massive, both is mass and volume.
Electronegativity Trends
• Period - electronegativity increases as you go from left
to right across a period.
• Why? Elements on the left of the period table have 1 -2
valence electrons and would rather give those few
valence electrons away (to achieve the octet in a lower
energy level) than grab another atom's electrons. As a
result, they have low electronegativity. Elements on the
right side of the period table only need a few electrons to
complete the octet, so they have strong desire to grab
another atom's electrons.
Electronegativity Trends cont.
• Group - electronegativity decreases as you go down a
group.
• Why? Elements near the top of the period table have few
electrons to begin with; every electron is a big deal. They
have a stronger desire to acquire more electrons.
Elements near the bottom of the chart have so many
electrons that loosing or acquiring an electron is not as
big a deal. This is due to the shielding effect where
electrons in lower energy levels shield the positive
charge of the nucleus from outer electrons resulting in
those outer electrons not being as tightly bound to the
atom.
Ionization Energy
• Ionization energy is the amount of
energy required to remove the
outermost electron/s.
• Ionization energy is closely
related to electronegativity.
Electronegativity
• Electronegativity is an
atom's 'desire' to grab
another atom's electrons.
Ionization Energy Trends
• Period - ionization energy increases as you go from left
to right across a period.
• Why? Elements on the right of the chart want to take
others atom's electron (not given them up) because they
are close to achieving the octet. The means it will require
more energy to remove the outer most electron.
Elements on the left of the chart would prefer to give up
their electrons so it is easy to remove them, requiring
less energy (low ionization energy).
25
Ionization Energy Trends cont.
• Group - ionization energy decreases as
you go down a group.
2) or how badly an atom wants
to take other atom's electrons
(electronegativity)
• Why? The shielding effect makes it easier
to remove the outer most electrons from
those atoms that have many electrons
(those near the bottom of the chart).
Reactivity of Metals
Period - reactivity decreases as you go from
left to right across a period.
Group - reactivity increases as you go down a
group
Why? The farther to the left and down the
periodic chart you go, the easier it is for
electrons to be given or taken away, resulting in
higher reactivity.
Reactivity
Reactivity refers to how likely or
vigorously an atom is to react with
other substances.
This is usually determined by two
things:
1) How easily electrons can be
removed (ionization energy)
from an atom
The transfer/interaction of
electrons is the basis of chemical
reactions.
26
Ionic Radius vs. Atomic Radius
Reactivity of Non-Metals
Period - reactivity increases as you go from
the left to the right across a period.
Group - reactivity decreases as you go down
the group.
Why? The farther right and up you go on the
periodic table, the higher the
electronegativity, resulting in a more
vigorous exchange of electron.
Ionic Radius vs. Atomic
Radius
Summary of Periodic Trends
• Metals - the atomic radius of a metal is generally
larger than the ionic radius of the same element.
• Why? Generally, metals loose electrons to achieve the
octet. This creates a larger positive charge in the
nucleus than the negative charge in the electron cloud,
causing the electron cloud to be drawn a little closer to
the nucleus as an ion.
Ionic Radius vs. Atomic Radius
cont.
Non-metals - the atomic radius of a nonmetal is generally smaller than the ionic
radius of the same element.

Why? Generally, non-metals loose
electrons to achieve the octet. This creates
a larger negative charge in the electron
cloud than positive charge in the nucleus,
causing the electron cloud to 'puff out' a little
bit as an ion.

2.8 Periodic Trends Activity Determining Periodic Trends Activity
Purpose
To understand the periodic nature of the periodic table – by knowing an element’s location on the
periodic table one can determine the element’s number of valence electrons and have an awareness of
its ionization energy, density, electron affinity, electronegativity, atomic radius, and ionic radius, relative
to other elements
27
Materials
Tiles from Periodic Trends, periodic table, scissors, tape, white paper, graph paper
Procedure
Part 1
1. You will be given cards for each of five elements in a certain group. Each card contains
information about the physical properties of the element. Use the descriptions in the Periodic
Trends section below to place your element cards in what you think is the correct arrangement.
You only have representative or main-group elements. The transition metals and noble gases
are not used in this activity.
2. Use a periodic table to predict the atomic number and period of each element and write them in
the spaces provided on the card.
Part 2
1. As a group, organize all the elements into a periodic table. Tape the elements onto white
paper when you are sure of their placement.
2. Compare your periodic table with the teacher’s key and make any necessary adjustments.
3. Complete the Analysis and Conclusions section.
Terminology of Periodic Trends
Valence electron: An electron in an outer shell of an atom that can participate in forming chemical
bonds with other atoms. Nonvalence electrons are tightly bound to the nucleus and are called core
electrons.
Octet rule: Atoms tend to gain, lose, or share electrons to reach eight electrons in their outer electron
shells.
Ionization energy: The amount of energy required to remove one electron from a neutral atom that is in
the gaseous state. Elements on the right side really want to keep their electrons because they are so
close to achieving an octet. Elements on the left side don’t mind losing an electron. Going down a group,
the shielding effect of additional electrons makes it easier to remove outermost electrons.
Atomic radius: One-half the distance of a single bond between two atoms of an element. As we go
down a group, the principal energy level increases, causing an increase in the average distance between
the electrons and the nucleus. Atoms get bigger as electrons are added to the principal energy levels. All
the atoms in a given period have their outermost electrons in the same principal energy level. As we go
across a period, we are adding electrons to the same principal energy level. The atoms do not get bigger
across a period because we are also adding protons to the nucleus as we move across a period. The
additional protons in the positive nucleus increase the pull on the negative electrons resulting in smaller
atoms across a period.
Ionic radius: One-half the center-to-center distance between two ions in a crystalline ionic compound.
Ions on the left side of the table are smaller than the atoms they come from because they’ve lost an
electron. Ions on the right side are bigger than the atoms they come from because they’ve gained an
electron.
Electronegativity: A measure of attraction of an atom for a pair of shared electrons, a measure of
whether an atom “grabs” more than its fair share of the shared electrons. As you go down a group,
atoms have more total electrons so they don’t really care that much about their outermost ones. As you
go across a period, atoms have more valence electrons and are so close to achieving an octet that they’ll
grab another atom’s electrons.
28
Analysis and Conclusions
1. Use the words increases or decreases to describe the trends for the following properties of the
representative elements, going from left to right across a period:
• ionization energy
• atomic radius
____________________________________
____________________________________
• ionic radius
____________________________________
• electronegativity
____________________________________
2. Use the words increases or decreases to describe the trends for the following properties of the
representative elements, going from top to bottom down a group:
• ionization energy
____________________________________
• atomic radius
____________________________________
• ionic radius
____________________________________
• electronegativity
____________________________________
3. When elements in Group I ionize, are they more likely to gain or lose electrons?
Explain your answer.
4. Explain why Group I elements have ionic radii smaller than their atomic radii.
5. When elements in Group VII ionize, are they more likely to gain or lose electrons?
Explain your answer.
6. Explain why Group VII elements have ionic radii larger than their atomic radii.
7. Compare how the melting point for metals and nonmetals generally changes as you move down a
group.
29
8. What is the relationship between number of valence electrons and group number?
9. Create a graph of ionization energy versus atomic number and identify the different periods. Compare
the ionization energies of metals to nonmetals.
10. Create a graph of atomic radius versus atomic number and identify the different periods. What is the
trend going down a group? What is the trend as you go across a period?
30
Unit 2: Homework
Hwk 2.1 Atomic Structure
Rutherford’s Gold Foil Experiment: view experiment in video online
Summarize the experiment and the major conclusions below:
_____________________________________________________________________________________
_____________________________________________________________________________________
_____________________________________________________________________________________
_____________________________________________________________________________________
________________________________
Atomic theory, as we know it today, is the result of the contributions of many scientists who did the research,
disproved old models and suggested the new ones, and added to the atomic theory. The items below are the
concepts that were developed in the history of the atomic theory. Rank them in order from oldest to most recent.
____Plum Pudding Model
____Electrons are in orbits
____Atoms of the same element have the same properties
____Positivity-charged atomic nucleus
Indicate whether each of the following statements about the nucleus of an atom is true or false.
_____The nucleus of an atom is neutral.
_____The nucleus of an atom contains only neutrons.
_____The number of nucleons present in the nucleus is equal to the number of electrons present outside
the nucleus.
______The nucleus accounts for almost all the volume of an atom.
______The nucleus accounts for almost all the mass of an atom.
______The nucleus can be positively or negatively charged depending on the identity of the atom.
31
Hwk 2.2 Ions, Isotopes and average atomic weight
Complete this table. Note that the atoms/ions are not necessarily neutral.
Element
Se
Mass Number
76
Number of Neutrons
65
36
39
Number of Protons
36
Number of electrons
36
Charge
-2
+1
Calculate the atomic mass of naturally occurring tungsten (W) to five sigfigs given the following isotopic masses
and abundances:
W=
______________________ amu
Isotope
Isotopic mass (amu)
Abundance )%)
𝟏𝟖𝟎
𝑾
179.946706
0.12
𝟏𝟖𝟐
𝑾
181.948206
26.50
𝟏𝟖𝟑
𝑾
182.9502245
14.31
𝟏𝟖𝟒
𝑾
183.9509326
30.64
𝟏𝟖𝟔
𝑾
185.954362
28.43
Suppose that a fictitious element, X, have two isotopes:
59
𝑋 (59.015 amu) and 62𝑋 (62.011 amu). The lighter
isotope has an abundance of 73.7%. Calculate the atomic mass of the element X.
Hwk 2.3 General Properties of Periodic Table
Identify these elements based on their location in the periodic table. Give the symbol and number
Period 5, group 14 (4A) __________________________________________________
Period 4, group 12 (2B)___________________________________________________
Period 5, group 18 (8A) ___________________________________________________
Label these groups on the periodic table below: Halogens, Alkali metals, Alkaline Earth metals, Noble
Gases, Transition Metals, Hydrogen, Helium, Metalloids, Non-metals, Lanthanide and Actinides. You can
use arrows of shade the blocks in color.
32
Classify each of the following elements as Metals (M), Metalloids (MO) and Non-metals (NM)
As, _______ Ga________, Tl_________, Xe_______, Si_______, S________, Bi_________
Hwk 2.4 Electron Conf & Quantum Numbers
Identify the atom with the following ground-state electron configuration for its valence shell.
2𝑠 2 2𝑝4
Symbol: ____________
Classify the following orbitals as s, p, d or f, according to their shapes:
What quantum numbers specify these subshells?
4s 2p 3d 4f
n=
l=
33
Hwk 2.5 More Electron Conf & QN
In the space below, write the unabbreviated electron configurations for the following elements:
1)
Barium
________________________________________________
2)
Neptunium
________________________________________________
In the space below, write the abbreviated electron configurations for the following elements:
1)
Radium
________________________________________________
2)
Lawrencium
________________________________________________
Determine what elements are denoted by the following electron configurations:
1.
[Xe] 6s24f145d6 ____________________
2.
[Rn] 7s25f11 ____________________
Determine which of the following electron configurations are not valid:
1. [Kr] 5s24d105p5 ____________________
1.
[Xe] ____________________
Hwk 2.6 Energy Calculations
1.
What is the energy of a photon with a wavelength of 6.55
102 nm?
2.
What is the wavelength of a photon with an energy of 3.40
10-18 J?
3.
What is the wavelength of a photon with a frequency of 2.35
1012 s
34
Hwk 2.7 Periodic Trends
Rewrite these elements from most to least electronegative. Al, Na, Rb, F, N ______________________
Rank these elements from largest to smallest radius. Ne, Li, B, N, F, C, O, Be _______________________
Rank these elements from highest to lowest ionization energy. Br, Kr, K, Ge, Ca, Se _________________
Rank these ions according to largest to smallest radius. N-3, F-, Mg2+,O2-, Na+ _______________________
Hwk 2.8 More Periodic Trends
1. What is the symbol for the following elements.
a. Magnesium _____________
b. Potassium ______________
2. What are the names of the following elements.
a. C __________________
b. Cl _________________
3. What period are the following elements in?
a. He _______________
b. Ge _________________
4. What group are the following elements?
a. Sulfur _______________
b. Ca _________________
5. Give me an atom with the following characteristics.
a. Halogen _________________
b. Nonmetal ________________
c. Alkali metal ______________
d. metalloid ________________
e. Lanthanide series __________
f. Alkaline Earth metal ________________
g. Transition metal ___________
h. Nobel gas ________________
6. Write the electron longhand and shorthand configurations and give the quantum #s for:
a. Li _______________________________________________________________
b. Na ______________________________________________________________
c. K _______________________________________________________________
7. What are valence electrons? __________________________________________________
8. How many valence electrons are in the following element?
a. F ________
b. Cl ___________
c. Br ____________
d. I _____________
e. O ________
f. S ___________
g. Se ____________
h. Te ____________
Hwk 2.9 Review
1. When compared to copper-63, copper-65 has more —
a. protons
b. neutrons
c. energy levels
configurations
2.
d. bonding
Which of the following pairs of atomic symbols represent isotopes of the same element?
a. 235U and 238U
b. P4 and P8
c. 32P and 83Pb
d. 50Sn and 51Sb
3. Which of the following represents a particle containing 8 protons, 9 neutrons and 8 electrons?
A. Oxygen-16
B. Nitrogen-17
C. Oxygen-17
D. Fluorine-15
35
4. Fill in the chart below:
Isotope
Isotope
Symbol
name
Remember:
Mass
number
Sodium-23
23
#of particles Number of
in the
protons
nucleus
23
Number of
neutrons
Number of
electrons
32
32
15𝑃
8
10
5. Define a family. _______________________________________________________
6. What is a period? ________________________________________________________
7. What is the symbol for the following elements.
a. Magnesium _____________
b. Potassium ______________
8. What are the names of the following elements.
a. C __________________
b. Cl _________________
9. What period are the following elements in?
a. He _______________
b. Ge _________________
10. What group are the following elements?
a. Sulfur _______________
b. Ca _________________
11. Give me an atom with the following characteristics.
a. Halogen _________________
b. Nonmetal ________________
c. Alkali metal ______________
d. metalloid ________________
e. Lanthanide series __________
f. Alkaline Earth metal ________________
g. Transition metal ___________
h. Nobel gas ________________
12. Write the electron longhand and shorthand configurations and give the quantum #s for:
a. Li _______________________________________________________________
b. Na ______________________________________________________________
c. K _______________________________________________________________
13. What are valence electrons? __________________________________________________
14. How many valence electrons are in the following element?
a. F ________
b. Cl ___________
c. Br ____________
d. I _____________
e. O ________
f. S ___________
g. Se ____________
h. Te ____________
On the blank periodic table below
15. Label the s, p, d, and f block elements
16. Create a circle that fills the whole box where the largest atom exists in the periodic table.
17. Put a dot where the smallest atom is in the periodic table.
18. Put a triangle on the box with the atom with the highest electronegativity
19. Put a square in the box with the lowest ionization energy
36
20. Label with arrows the trends for: atomic radius, ionic radius (metals and nonmetals), ionization
energy, and electronegativity
From Electron Configuration Notes:
17. What is the shape of the s orbital? _________p orbital?___________ d orbital? _________________
18. Which of the following orbitals is closest to the nucleus?
a. 2s
b. 3p
c. 1s
d. 4d
19. In the wave-mechanical (quantum) model of the atom, orbitals are regions of the most probable
locations of:
a. protons
b. positrons
c. neutrons
d. electrons
20. (CHALLENGE) Heiseinberg’s Uncertainty Principle states:
_____________________________________________________________________________________
___________________________________________________________________
21. Identify the following atom, 1s22s22p63s23p64s23d6 _________________________
22. Write out the orbital notation and give the quantum numbers for F, N, Zn:
23. Write the electron configuration for the above elements:
24.Write the shorthand (noble gas) notation for the above elements:
F:
__________________________________________
F1— :
__________________________________________
Zn:
__________________________________________
Zn2--:
_________________________________________
Sr:
_________________________________________
Sr1+:
_________________________________________