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Chapter 5: The Periodic Law. CP Chemistry Mrs. Klingaman. #13) Valence Electrons. Packet pg. 4. Definition : The electrons available to be lost, gained or shared in the formation of chemical compounds The OUTERMOST (highest energy level) electrons in an atom

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chapter 5 the periodic law

Chapter 5: The Periodic Law

CP Chemistry

Mrs. Klingaman

13 valence electrons
#13) Valence Electrons

Packet pg. 4

  • Definition:
    • The electrons available to be lost, gained or shared in the formation of chemical compounds
    • The OUTERMOST (highest energy level) electrons in an atom
  • For main-group elements, the valence electrons are the electrons in the highest energy level s & p atomic orbitals
3 valence electrons
#3) Valence Electrons

Packet pg. 10

  • Definition:
    • The OUTERMOST (highest energy level) electrons in an atom
    • Electrons involved in the chemical bonding of atoms
5 bohr diagrams
#5) Bohr Diagrams

Packet pg. 10

  • Show total number of electrons in an atom and to which energy levels those electrons belong
    • Drawn like the Bohr Model, where each ring is an energy level (“n”)
    • Electrons fill in the lowest energy levels first (closest to nucleus)
5 bohr diagrams1
#5) Bohr Diagrams

Packet pg. 10

  • Each energy level has a different maximum number of electrons which it can contain
5 bohr diagrams2
#5) Bohr Diagrams

Packet pg. 10

  • Nucleus:
    • Determine # of protons & neutrons
  • Energy level rings:
    • The total number of energy levels in an atom is the same as the period # the atom is in
    • Fill in all electrons on the appropriate energy level (keep in mind the maximum # of electrons allowed on each level
oxygen bohr diagram
Oxygen - Bohr Diagram
  • p+ _______; n0 _______;
  • e– _______
  • Group #: ____________
  • Period #: ____________
  • How many valence electrons does oxygen have? _____________

Hint: Look at the outer most ring, this is the HIGHEST energy level (the 2nd level)

silicon bohr diagram
Silicon- Bohr Diagram

Silicon (Si) : 1s22s22p63s23p2

  • p+ _______; n0 _______;
  • e– _______
  • Group #: ____________
  • Period #: ____________
  • How many valence electrons does oxygen have? _____________

Hint: Look at the outer most ring, this is the HIGHEST energy level (the 2nd level)

6 lewis dot structures
#6) Lewis Dot Structures

Packet pg. 10

  • Show only the valence electrons in an atom (represented by dots)
  • No atom can have more than 8 valence electrons

X

3

7

Dots (# of valence electrons) are added around the element symbol so that one electron is on each side first

6

5

1

2

8

4

lewis diagrams
Lewis Diagrams:
  • Examples
    • Na – Cl
    • C – Ne
    • S
slide16

Bohr & Lewis - Practice

Chlorine

  • p+ _______; n0_______; e– _______
  • Group Name: __________________________
  • Group #: ________ Period #: ________
  • Draw Bohr Diagram & Lewis Dot Structure
chapter 5 periodic trends

Chapter 5: Periodic Trends

CP Chemistry

Mrs. Klingaman

periodic trends
Periodic Trends
  • Group Trends
    • Describes the general trend for elements looking down a group/family
    • from top to bottom
  • Period Trends
    • Describes the general trend for elements looking across a period
    • from left to right
periodic trends1
Periodic Trends
  • There are 4 periodic trends we will discuss:
    • Atomic Radius
    • Ionization Energy
    • Electron Affinity
    • Electronegativity
7 atomic radius
#7) Atomic Radius

Packet pg. 2

  • Definition:
    • ½ the distance between the nuclei of identical atoms bonded together
    • The relative size of an atom
7 atomic radius ar
#7) Atomic Radius (AR)

Packet pg. 2

  • Group Trend:
    • generally increases down a group.
    • looking top-to-bottom down a group, atoms get larger in size (measured by atomic radius)
  • Why?- Increasing Energy Level
    • As you move down a column, each successive element has one additional energy level surrounding the nucleus of the atom
    • This means electrons in the atom are farther away from the nucleus, creating a larger atomic radius
slide22
Atomic RadiusGroup Trend:Increasesdown a column due to increasing energy level (electrons are farther from nucleus)
7 atomic radius ar1
#7) Atomic Radius (AR)

Packet pg. 2

  • Period Trend:
    • generally decreasesacross a period.
    • looking left-to-right across a period, atoms get smaller in size (measured by atomic radius)
  • Why?- Increasing Nuclear Charge
    • As you move across a period, each successive element has one more proton in its nucleus than the previous atom.
    • This increasing positive nuclear charge will cause a greater force of attraction and pull the electrons in closer to the nucleus.
slide24
Atomic RadiusPeriod Trend:Decreases across a period due to increasing nuclear charge(electrons are pulled closer to the nucleus)
slide25

Atomic Radius: Hint Element

Francium (Fr) = “Frankenstein”  the biggest

Fluorine (F)  smallest

skip to 10 ions
Skip to #10) Ions

Packet pg. 3

  • What are Ions?
    • An atom or group of bonded atoms that have a positive or negative CHARGE
    • Unlike atoms which are electrically NEUTRAL, ions have a net CHARGE (either positive or negative)
  • How are Ions Formed?
    • Formed by either adding or removing electrons to/from a neutral atom
    • Adding electronswill form a NEGATIVE ion
    • Removing electronswill form a POSITIVE ion
11 cation
#11) Cation

Packet pg. 3

Cation:

  • Positively charged ion
  • Formed by the loss on 1 or more electrons from a neutral atom
  • On the periodic table, metals on the left side form cations
12 anion
#12) Anion

Packet pg. 3

Anion:

  • Negatively charged ion
  • Formed by the addition of 1 or more electrons from a neutral atom
  • On the periodic table, nonmetals on the right side form anions
skip to 15 ions valence electrons
Skip to #15) Ions & Valence Electrons

Packet pg. 4

  • The maximum number of valence electrons an atom may have is ______. When this occurs the atom is said to have a full valence shell.
  • The elements within group _______ (_____________) do have a full valence shell; This full valence shell is the reason why elements in this group are stable and chemically unreactive.

8

18/8

Noble Gases

skip to 15 ions valence electrons1
Skip to #15) Ions & Valence Electrons

Packet pg. 4

  • In fact, all chemically bonding occurs because every element wants to be stable and have a full valence shell. In order to achieve this, some atoms will lose electrons (to form ________) and some atoms will gain electrons ( to form _______)

cations

anions

8 ionization energy ie
#8) Ionization Energy (IE)

Packet pg. 2

  • Definition:
    • the energy required to REMOVE an electron from a neutral atom of an element
    • Equation to represent this:

A + energy  A+ + e–

Atom + energy  ion w/positive charge + electron that was

aka: “cation” removed

8 ionization energy ie1
#8) Ionization Energy (IE)

Packet pg. 2

  • Group Trend:
    • generally DECREASESdown a group.
    • it takes LESS energy to remove electrons from the atoms farther down in a group
  • Why?- Increasing Energy Level
    • As you move down a column, each successive element is larger and has electrons in higher energy levels so the electrons are farther away from the nucleus
    • Electrons farther away from the nucleus, are not as greatly attracted/pulled in by the nucleus, thus LESS energy is required to rip an electron away from the bigger atoms
8 ionization energy ie2
#8) Ionization Energy (IE)

Packet pg. 2

  • Period Trend:
    • generally INCREASESacross a period.
    • it takes MORE energy to remove electrons from the atoms across a period
  • Why?- Increasing Nuclear Charge
    • Again we see that moving across a period, each successive element has one more proton in its nucleus than the previous atom.
    • This increasing positive nuclear charge will cause a greater force of attraction and pull the electrons in closer to the nucleus, making it HARDER to remove the electrons and thus requiring MORE energy
slide34

Ionization Energy: Hint Element

Helium (He) the HIGHEST ionization energy

b/c it takes the most energy to remove electrons

In general, the Noble Gas Family has the HIGHEST ionization energies of any family.

Why?

9 electron affinity ea
#9) Electron Affinity (EA)

Packet pg. 3

  • Definition:
    • the energy associated with the addition of an electron to a neutral atom
    • Equation to represent this:

A + e– A– + energy

Atom + add an electron ion w/negative charge + energy

aka: “anion” released

9 electron affinity ea1
#9) Electron Affinity (EA)

Packet pg. 3

  • Difference between EA and IE:
    • EA is basically the opposite of IE
    • Electron Affinity = energy is released when an electron is ADDED to an atom
    • Ionization Energy = energy is requiredwhen an electron is REMOVED from an atom
9 electron affinity ea2
#9) Electron Affinity (EA)

Packet pg. 3

  • Group Trend:
    • generally DECREASESdown a group.
    • Down a column, electrons added to atoms will have LESS energy released than the atoms above them
  • Why?- Increasing Energy Level
    • Down a group, atoms get larger with more energy levels; larger atoms contain more electrons inside them
    • The more electrons there are in the atom, the LESS likely the nucleus will be to pull/attractadditional electrons into the atom; when finally an electron is gained by the atom, LESS energy is released because it was harder to gain that electron
9 electron affinity ea3
#9) Electron Affinity (EA)

Packet pg. 3

  • Period Trend:
    • generally INCREASESacross a period.
    • Across a period, MOREenergy is released when electrons are added
  • Why?- Increasing Nuclear Charge
    • Again we see that moving across a period, each successive element has one more proton in its nucleus than the previous atom.
    • This increasing positive nuclear charge will cause a greater force of attraction and pull the electrons in closer to the nucleus, making itEASIER to gain an electron, thus releasing MORE energy when the electron is added to the atom
slide39

Electron Affinity: Hint Element

Fluorine (F) the MOST energy released when an electron is added to an atom

The Halogens have the highest EA because it is really beneficial if they gain an electron  = lots of energy released

16 electronegativity en
#16) Electronegativity (EN)

Packet pg. 5

  • Definition:
    • the ability of an atom in a chemical compound to attract or pull electrons to itselffrom another atom
16 electronegativity en1
#16) Electronegativity (EN)

Packet pg. 5

  • Group Trend:
    • generally DECREASESdown a group.
    • Going down a column, atoms are LESS able to attract electrons towards itself and away from other atoms in a compound
  • Why?- Increasing Energy Level
    • Down a column, atoms are larger with more & more electrons surrounding the nucleus
    • The more electrons surrounding the nucleus, the LESS of an attractive force the nucleus will have in order to pull electrons away from other atoms.
16 electronegativity en2
#16) Electronegativity (EN)

Packet pg. 5

  • Period Trend:
    • generally INCREASESacross a period.
    • Across a period, atoms have a GREATER ability to attract electrons to itself
  • Why?- Increasing Nuclear Charge
    • Again we see that moving across a period, each successive element has one more proton in its nucleus than the previous atom.
    • This increasing positive nuclear charge will cause a stronger force of attraction between the nucleus and electrons (even electrons in other atoms)
    • The higher the nuclear charge, the GREATER the pull the atom will have in order to attract electrons towards itself
16 electrongeativity en
#16) Electrongeativity (EN)

Packet pg. 5

  • To rank atoms based on this ability, each element is given an EN value between 0.0 and 4.0 (the Linus Pauling EN Scale)
  • The higher the value (closer to 4.0) the more EN the atom
  • The Noble Gases are not even listed or given EN values because they have ZERO ability to attract electrons toward themselves!!!
linus pauling en scale

Electronegativity: Hint Element

Fluorine (F) the GREATEST ability to attract electrons to itself (very tiny w/a high nuclear charge)

Linus Pauling EN Scale
electron affinity and electronegativity
Electron Affinity and Electronegativity
  • EA & EN are similar to one another:
    • Both have the same group & period trends
    • Both deal with gaining or attracting electrons
  • Difference between EA and IE:
    • Electron Affinity = the ENERGY CHANGE associated with gaining and electron
    • Electronegativity = the ABILITY, or how well, and atom can attract an electron