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UNIT 4
Learning Objectives For Electron Configuartion & Periodic Table
Students are to demonstrate the following through a Unit TEST or a LAYERED ASSIGNMENT.
Unit TEST on :_________
Layered Assignment Due:_________
1.
Review atomic structure: definition, location and characteristics of the three major subatomic
particles (the proton, the neutron and the electron).
2.
Use the periodic table and its key to determine the symbol, name, atomic number, relative
average atomic mass, number of protons and number of electrons (in a neutrally charged atom)
of an element.
3.
Describe atomic orbitals in terms of their shape, relative size, and relative energy.
4.
Locate electrons in energy levels, sublevels and orbitals using principles of orbital energy, orbital
capacity and electron spin.
5.
Make orbital diagrams and write electron configurations of atoms, ions and isoelectric species.
6.
Define the term: periodic trend and use periodic trends to predict relative characteristics of
elements
7.
Name the families of elements and identify the properties related to each family.
8.
Describe trends in the properties and behavior of elements as you proceed through the periodic
table.
9.
Explain why elements in the same group/family have similar properties: the relationship between
family placement and electronic configuration.
PBA LAB REPORT Due:____________ Characteristic Flame Test Experiment Performed on: __________
Textbook (Red) Homwork Due:________
Set 1 DO on:_____________ p129 #62, 73 p166 #38, 60, 64, p167 #78, 81 Atomic Structure
Set 2 DO on:_____________ p 167 #70, 71, 82, 83, 84, 85, 86 p168 # 87, 90 p 199 # 69 p200 #88 Electron Configuration
Set 3 DO on: ____________ p 198 #28, 29, 31, 32, 42, 44, 46, p199 #48, Periodic Table
Set 4 DO on: ____________ p 198 #58, 60, 63, p 200 #78, 81, Periodic Trends
SEE POWER SCHOOL for Formative Assessments/ Practice / Activities/ Date Changes
Physical and chemical properties are a function of electron configuration because they are directly related to
‘coulombic attraction”: quantity of charge and distance separating charge.
Electron Configuration of outer most shell : valence shell are the same with in a group/family.
Same outer most/valence shell configuration different energy levels with in a group/family
Lewis Dot Diagram shows only the element symbol and the VALENCE ELECTRONS.
This is WHY elements periodically repeat their chemical properties by their family/group.
Isoelectric Species: Have same electron configuration.
Many elements form ions to become isoelectric with nearest noble gas (circular periodic table)
Trends on periodic table: be able to explain
Ionization energy: energy needed to remove an electron from atom, an actual measured value
Ionization energy
decreases down a column b/c larger atom electrons are farther from nucleus, less attraction
increases across a row b/c larger # of protons in nucleus at same distance/shell, more attraction
Atomic radius: size of an atom, an actual measured value
Atomic radius
increases down a column b/c larger atom electrons are farther from nucleus
decreases across a row b/c larger # of protons in nucleus at same distance/shell, more attraction
Electronegativity: a calculated value to express ability of an atom to attract electrons itself in a bond (typical covalent)
Electronegativity:
decreases down a column b/c larger atom electrons are farther from nucleus, less attraction
increases across a row b/c larger # of protons in nucleus at same distance/shell, more attraction
Metallic Character: meaning luster, ability to be deformed (malleable/ductile), conductors of heat and electricity: electrons
move easily
increases down a column b/c low electron attraction: electrons free “float in a sea of electrons”
decreases across a row b/c more attraction: held tightly
Atomic structure of elements:
Atomic number = number of protons
Mass number = number of protons plus
neutrons
In a neutral atom: # protons = # electron.
Charged ion: # protons stays the same.
More electrons if negative, Less electrons if positive.
Protons are positive
Neutrons are neutral
Electrons are negative
Protons count as mass of 1 amu
Neutrons count as mass of 1 amu
Electrons so small essentially zero mass
Atomic mass on periodic table is a relative average mass on periodic table is average of other whole number mass numbers.
# of protons dictates type of element, regardless of how many or how few neutrons or electrons.
# protons plus #neutrons = mass number (atomic mass is average of all isotopes
# protons minus #electrons = net charge (neutral/zero charge called atom, ions: positive chargecation, negative chargeanion)
Chemical symbols are abbreviations for the names of the elements. Only one capital letter.
Many element names come from old latin and greek words e.g. Na is sodium (In latin it is natrium)
The periodic table is a list of all the elements that is known today. It is a constantly increasing list.
Arranged so that metals are on the left side, nonmetals on the right and metalloids along the zig zag line.
Period is a horizontal row, there are 7 periods
Group / Family is a vertical column (some times a block)
Periodic law: Physical and chemical properties are periodic functions of atomic number
Doeberiner: triads of similar element/Newlands : law of octaves/Meyer: closer to modern
Mendeleev: credited with first periodic table, predicted missing elements and their properties!
Periodic law: Physical and chemical properties are periodic functions of atomic number
Metals are usually solids at room temperature except for mercury, shiny, can be bent and beaten flat without breaking
malleable/ductile, have high melting points and are good conductors of heat and electricity.
Nonmetals can be solids, liquids or gases. Are dull in color. Have low melting points, are brittle, and are poor conductors of heat
and electricity.
Metalloids are combination of the two.
Periodic Table 7 rows called periods: 18 columns called groups/families
Families of elements:
Type of ion formed
Noble gases: column 18
none; consider inert (non reactive/doesn’t bond)
Halogens: column 17
-1 (nonmetalanion, one column less than noble gas)
Chalcogens/Oxygen: column 16
-2( nonmetalanion, two column less than noble gas)
Alkali Metals: column 1
+1(metalcation, one column more than noble gas, circular per.table)
Alkaline Earth Metals: column 2
+2(metalcation, two columns more than noble gas,circular per. table)
Transition Elements: columns 3-12
Multiple ions may form
(Inner Transition Elements: Lanthanide & Actinide series: Most radioactive
Hydrogen: acts both like a halogen and like an alkali metal, considered family by itself, forms +1 and sometimes –1
Be able to write the chemical formula and chemical name for ionic compounds studied so far and be able to calculate their
corresponding molar mass. Remember criss-cross, simplify, “poly wants parentheses.”
Cations—atoms that have lost electrons, carry a positive charge.
Group I Alkali Metals
+1
Group II Alkaline Earth Metals
+2
Group III (Al)
+3
Anions—atoms that have gained electrons, carry a negative charge.
Group VI (Chalcogens)
-2
Group VII (Halogens)
-1
Group V (N, P,)
-3
Electron Configuration of outer most shell : valence shell are the same with in a group/family.
Same outer most/valence shell configuration different energy levels with in a group/family
Lewis Dot Diagram shows only the element symbol and the VALENCE ELECTRONS.
Bohr Model: “energy shells” replaced by Quantum mechanical model
Explain how Bohr’s model of the atom incorporated Plank’s idea of quantization.
The difference between a line spectrum and a continuous spectrum was the evidence for the Bohr Model and later, when
electrons were considered to have a dual wave – particle character, for the Quantum Mechanical Model..
The Heisenberg’s Uncertainty Principle and explains that it is impossible to know both the precise location and precise
momentum (mass x velocity) of an electron at the same time.
.
Therefore, the location of an electron is discussed as a “probability” and it shown by electron density. The 90% probability line is
used as an “outline” for the shape of an “orbital”
The s,p,d,and f orbitals in terms of shape, size and energy are mathematically equivalent to the bright line spectrums
(Schrödinger’s Wave Equation).
Recognize, and draw the shapes of the s and p orbitals (s = spherical, p = dumb bell) (NOT ORBITS). Orbitals are three
dimensional probability maps of the electron wave. They come from solving Schrodinger’s wave function for the hydrogen atom
State the Pauli Exclusion Principle
The four quantum numbers: n, l, ml, and ms:
ENERGY LEVEL
Principle quantum number, n. –Determines the energy of the electron in a given shell and the size of the atomic orbital.
SUBLEVEL
Angular quantum number, l.
-Describes the shape of the electron orbital also called sublevels.l = (0, 1, 2, 3, 4, …n-1)
ORBITAL
Magnetic quantum number, ml. -Describes the directionality of the electron orbital.
SPIN
Spin quantum number, ms
Ml = (-l, … +l)
-Described as +1/2 or -1/2 to indicate direction of spin.
The electron configuration of an element using the Pauli Exclusion Principle, the Aufbau principle, and Hund’s rule are meant to
describe the electrons’ location.
Follow the general filling (long hand) trend of electrons into atomic orbitals using the following order:
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 …
OVERLAP of evergy levels and sublevels“d” is filled one behind “s,” and “f” is filled two behind the “s.”
Write electron configurations using box/orbital diagrams where each electron is represented as a half arrow
Make the “d” orbital half filled when you are one electron away. Cr and Mo are written as follows [Ar] 4s 1 3d5 and [Kr] 5s1 4d5.
Note the promotion of the “s” electron to make a half filled “d” orbital.
Make the “d” orbital completely filled when you are one electron away. Cu and Ag are written as follows [Ar] 4s 1 3d10 and [Ag]
5s1 4d10
Write electron configurations for cations and anions.
When forming an ion, the valence shell electrons are removed first. Ca2+ [Ar]. The 4s2 electrons are removed first.
For anions, the electrons are added to outer, unfilled orbital.
Interpret the noble gas short hand for writing electron configurations. [Ne] = 1s2 2s2 2p6
Periodic table divided into s, p, d and f blocks and the periodic table maps with atomic orbitals:
Group I-II (s block) – Maximum 2 electrons
Group III-VIII (p-block) – Maximum 6 electrons
Transition metals (d-block) - Maximum 10 electrons
Lanthanides/actinides (f-block) - Maximum 14 electrons
Principle Energy Level (Like Bohr’s Shell) n = 1, 2, 3,4, 5, 6, 7…
Sublevel: defines shape of orbitals:
n=1 has 1 sublevel : 1s
n=2 has 2 sublevels: 2s, 2p
n=3 has 3 sublevels: 3s, 3p, 3d
n=4 has 4 sublevels: 4s, 4p, 4d, 4f
n=5 has 5 sublevels: 5s, 5p, 5d, 5f, 5 yet to be named
n=6 has 6 sublevels: 6s, 6p, 6d, 6f, 6 yet to be named, 6 also yet to be named
n=7 has 7 sublevel: 7s, 7p, 7d, 7f, 7yet to be named, 7 also yet to be named, 7 you get the idea
n=8 not yet discovered
Isoelectric Species: Have same electron configuration.
Many elements form ions to become isoelectric with nearest noble gas (circular periodic table)
Trends on periodic table: be able to explain
Ionization energy: energy needed to remove an electron from atom, an actual measured value
Ionization energy decreases down a column b/c larger atom electrons are farther from nucleus, less attraction
increases across a row b/c larger # of protons in nucleus at same distance/shell, more attraction
Atomic radius: size of an atom, an actual measured value
Atomic radius
increases down a column b/c larger atom electrons are farther from nucleus
decreases across a row b/c larger # of protons in nucleus at same distance/shell, more attraction
Electronegativity: a calculated value to express ability of an atom to attract electrons itself in a bond (typical covalent)
Electronegativity:
decreases down a column b/c larger atom electrons are farther from nucleus, less attraction
increases across a row b/c larger # of protons in nucleus at same distance/shell, more attraction
Metallic Character
increases down a column b/c low electron attraction: electrons free “float in a sea of electrons”
meaning more luster, easily deformed, good conductors electrons move
decreases across a row b/c more attraction: held tightly
Study Suggestions:
Chapter summaries
Class notes
Textbook reading
In Class / Practice worksheets
Homework Assignments
Key terms:
Quantum mechanics-highly mathematical theory of atomic structure based on the belief that energy is absorbed and radiated in
definite units called quanta since the energy of the electron is restricted to definite levels.
Principal quantum number-the number of the electron shell or principal energy level in an atom. It is represented by the symbol,
n. Determines the energy of the electron in a given shell. It also determines the size of the atomic orbital. n = (1, 2, 3, …
infinity)only integers and never zero.
Orbital-that part of the atom where the electron wave has a high probability of being found. Sometimes referred to as a sublevel.
This is a feature of the quantum mechanical model of the atom proposed by Schrodinger.
Heisenberg’s Uncertainty Principle-states that it is impossible to know both the precise location and precise momentum (mass x
velocity) of an electron at the same time.
Valence electrons-the electrons in the outermost shell of an atom.
Electron configuration-a notation describing the distribution of electrons among the sublevels or orbitals of the atom.
Aufbau principle—fill the lowest energy orbital first. “BOTTOMS UP” follows “n” as your guide, but remember that “d”
orbitals fill one behind the “s” and “f” orbitals fill two behind the “s
Pauli Exclusion Principle—no two electrons in an atom can have the same four quantum numbers. Make electrons spin paired
Hund’s Rule—“empty bus seat rule”. When several orbitals of equal energy are available, as in the p or d sublevels, electrons
enter singly with the maximum unpaired electrons.
Principle quantum number, n. –Determines the energy of the electron in a given shell. It also determines the size of the atomic
orbital. n = (1, 2, 3, … infinity)
only integers and never zero.
Angular quantum number, l. Describes the shape of the electron orbital also called sublevels. l = (0, 1, 2, 3, 4, …n-1)
Magnetic quantum number, ml. Describes the directionality of the electron orbital. M l = {-l,…0,…+l}
Spin quantum number, ms = +1/2 or -1/2
spin is physical property analogous in the macroscopic world to angular
momentum. The “turning electron” produces a vector force perpendicular to the direction of rotation (left hand rule).
Ground State-when all electrons in an atom are in the lowest energy levels available.
Excited State-when some of the electrons in the atom are at higher energy levels than the ones they normally occupy.
Line Spectra are empirical evidence used by Bohr to suggest that the energy of the electron in the atom is quantized in the atom.
Photoelectric effect-ability of light of certain frequency to eject electrons from the surface of a metal. It was explained by
Einstein using Planck’s concept of energy quanta.
Paramagnetic-the magnetic attraction of atoms, ions, or molecules having unpaired electrons.
Matter waves- De Broglie’s proposal that all matter has a dual wave-like character.
Photons—quanta (packets) of light energy
Electromagnetic spectrum consists of gamma rays, X-rays, UV, Visible, IR, microwaves, TV waves, and radio waves. Gamma
rays have the highest energy ,shortest wavelength and highest frequency while radio waves have the lowest energy, longest
wavelength and lowest frequency. Explain which color of light has more energy and why
Visible light has a wavelength range from 400 nm to 700 nm. Red has longer wave length than violet.