Welcome to
SCH3U!
Lesson 3
Structure and
Properties of
Elements
Learning goals and success criteria
Learning Goals
We are learning to
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explain the organization of the periodic table as well as periodic law
use an element’s position on the periodic table to explain its chemical reactivity, atomic
radius, ionization energy, and electron affinity
use the Bohr-Rutherford atomic model to explain trends in these properties
Success Criteria
I am able to
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describe the location of the electrons of the first 20 elements in terms of energy levels
explain and identify valence electrons of atoms
locate element families on the periodic table
describe the patterns in electron arrangement as one moves from top to bottom or from
left to right on the periodic table
explain the reasons for periodic trends in atomic size, ionization energy, and
electronegativity
use experimental observations of metal reactivity to create an activity series
Minds on: Reactivity of Alkali Metals
Beaker 1 (Left) - Lithium (Li) and water
Beaker 2 (Middle) - Sodium (Na) and Water
Beaker 3 (Right) - Potassium (K) and water
After you watch...
● What do you think it is about atoms of lithium, sodium, and potassium that cause
them to behave in such different ways?
● Hint: Draw the Bohr-Rutherford diagrams for the three elements and use those
drawings to form a rationale for reactivity.
● We will revisit this experiment at the end of the learning activity
Bohr Rutherford Model
● This model is a simplistic representation of the atom
○ the Nucleus with the protons and neutrons is shown in the centre
○ the electrons surround the nucleus on orbits (circles) called energy levels
Activity: Build an Atom Simulator
https://courseware-openhouse.ilc.org/sch3u_html/ilo/build_an_atom_phet/index.html
Compare Helium (He) and Beryllium (Be)
Maximum electron in each energy level
in the first 20 elements (2, 8, 8, 2…)
level
Maximum number of electrons
(true for the first 20 elements)
First
2
Second
8
Third
8
Fourth
2
Diagrams
Bohr-Rutherford diagrams
● In Bohr-Rutherford Diagrams, energy level 1 is located closest to the
nucleus, and each higher level is farther and farther away from
nucleus.
● Note: “When filling in the electrons on a Bohr-Rutherford diagram, each of the lone (unpaired) electrons are placed by themselves
first as seen in the diagrams above.”
Try this!
Draw Bohr-Rutherford diagrams for carbon-14 and Aluminum-27
Carbon -14
Aluminum-27
Valence electrons
The Outer energy level is referred to as the valence level , and the
electrons in the valence shell are called valence electrons.
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valence electrons determines how the periodic table of elements
is organized
Try this!
Name all the elements of the first 20 from the periodic table that have
four valence electrons
The periodic table
The periodic table was first developed by a Russian chemist named
Dmitri Mendeleev.
● This is what Mendeleev’s periodic table looked like:
● Arranged by Atomic mass
Watch the video The periodic Table: Crash Course
Mendeleev’s history
The genius of Mendeleev’s periodic table
Today’s periodic table
● Today, the periodic table is arranged by atomic number, not
atomic mass
○ when the elements are arranged by increasing atomic number,
it shows a periodic repetition in their properties
○ periodic law - the properties of elements that change as you
go across the table repeat themselves when you come to the
next row.
Features of the Periodic Table
● Periodic table is arranged in rows and columns.
● Elements are grouped like instruments in an orchestra
○ In an orchestra, loud instruments are at the back and quieter ones
at the front, and similar instruments grouped together
○ In the periodic table, larger elements are at the bottom, smaller
ones at the top, and similar elements are grouped together
Orchestra
Rows
● Rows are called periods and are numbered 1 through 7.
Columns
● Columns are called families or groups of elements, and are
both with the A and B and the numbers 1 to 18.
numbered
Metals vs. Non-metals
Bohr-Rutherford diagrams and the periodic table
Examine the following image. What patterns do you notice?
Try This!
1.
Describe the pattern of the number of electrons in the outer energy
level across a period.
#1 Answer
As you go across a period, the outer energy level gains one more
electron with each element.
Try This!
2. For all the A families, describe the relationship between the family
number and the number of electrons in the outer energy level.
#2 Answer
The A family number is the same as the number of electrons in the
outer level.
For example, family IA, all the elements have one electron in the outer
energy level.
Try This!
3. Describe similarities and differences in the number of electrons in the
outer energy level within a family (group).
#3 Answer
As you go down a group or column, the outer energy level has the same
number of electrons as the previous element. the only difference is that
a new energy level has been formed.
Periodic Trends
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Scientists have observed three of the “periodic trends”
○ Atomic radius, ionization energy, and electronegativity
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The periodic trends are related to the arrangement of protons and
electrons in an atoms
Atomic structure and periodic trends
Atomic radius - the distance from the centre of the nucleus to the
outermost electron of an atom
Trend in Atomic radius
● Across a period (from Left to right)
○ Atomic radius decreases (atomic size gets
smaller)
○ the number of protons increases = atomic
number
○ greater attraction between the valence
electrons and the nucleus
Trend in Atomic radius
● Down a group (from Top to Bottom)
○ Atomic radius increases (atomic size gets
bigger)
○ the number energy level (shells) increases
○ less attraction between the valence
electrons and the nucleus
Trend in Atomic radius
First Ionization Energy
● Some atoms can lose electrons and become ions (ionization) with an
overall + charge.
● First ionization energy (I.E.) - energy required to break the force of
attraction and remove the outermost electron from an atom
○ Two factors:
■ distance: the close the nucleus is to the nucleus, the stronger
attraction will be, and thus the greater ionization energy will be
needed to break that attraction
■ number of protons: the greater number of protons an atom has,
the more ionization energy needed to remove the electron from
an atom.
Trend in First Ionization Energy
● Across a period (Left to right)
○ ionization energy increases
why?
● number of protons increases
● attraction between valence
electrons and nucleus
increases
● the atom need more energy
to remove the electron from
the valence shell
● Down a group (Top to bottom)
○ ionization energy decreases
why?
● more energy levels (shells)
● valence electrons get
farther away from the
nucleus
● need less energy to remove
an electron
First Ionization Energy Trend
Electronegativity
● Electronegativity - the ability of an atom to attract a bonding
electron to itself
● same factors as the ionization energy: distance and the number of
protons
Electronegativity Trend
● Across a period (Left to right)
○ electronegativity increases
● Down a group (Top to bottom)
○ electronegativity decreases
why?
why?
● number of protons increases
● attraction between valence
electrons and nucleus
increases
● higher attraction for electrons
● more energy levels (shells)
● valence electrons get farther
away from the nucleus
● lower attraction for
electrons
● In general, metal elements have a weak attraction for electrons, while non-metals have
a strong attraction. Thus, the non-metals have a higher electronegativity than metals.
Electronegativity Trend
Consolidation - Exit ticket
Revisit the Minds on Activity on “Reactivity of Metals”
Answer the following questions
1. Which metal was the most reactive? (Na, Li, K)
2. What periodic trend is most important for determining metal
reactivity?
3. Which one is more reactive? rubidium (Rb) or cesium (Cs)?
Metals Reactivity Answers
1. Metals Reactivity in increasing order:
Li < Na < K
1. electronegativity
Because metals want to remove their outer
electron (lower attraction for electrons is
required)
K+H2O→KOH+H2(g)
1. Cesium is more reactive than Rubidium
because cesium has more energy levels, hence
the distance from electrons and protons are
farther away. This makes cesium easier to lose
electrons since the attraction between