Review - Element Properties

1
Review - Element Properties

• Critical Atomic Properties
• Electron Configuration
• Atomic Size
• Ionization Energy
• Electronegativity

2
Review - Element Properties

• Electron Configuration
• (nl) The distribution of electrons in the
  energy levels and sublevels of an atom
• n principal quantum indicating the energy and
  distance of orbitals from the nucleus
• Higher value of n means higher energy and
  greater distance from nucleus
• l Angular Momentum
• indicates shape of orbitals
• Values l 1 ? n -1
• ml Magnetic Quantum
• Values -l . 0 ..l

ml 0 (s orbital ml -1 0 1 (p orbitals) ml
-2 -1 0 1 2 (d orbitals) ml -3 -2 -1 0
1 2 3 (f orbitals)
3
Review - Element Properties

• Electron Configuration
• Quantum Numbers

4
Review - Element Properties
5
Review - Element Properties

• Electron Configuration
• Ground state Lowest Energy Level
• s block elements (groups 1 2)
• p block elements (groups 3, 4, 5, 6, 7, 8)
• Outer electron configurations (valence electrons)
  are similar within a Group
• Outer electron configurations (valence electrons)
  are different within a Period
• Outer electrons occupy the ns and np
  sublevels
• Four valence-level orbitals (one ns and 3 np)
  occur among the main-group elements
• The A group number (1A 8A) equals the number
  of valence electrons

6
Review - Element Properties

• Electron Configuration
• Outer electrons are shielded from the full
  nuclear charge by electrons in the same level and
  even more so by the electrons in the lower levels
• Shielding reduces attraction of nuclear charge
  resulting in a reduced effective nuclear charge,
  Zeff)
• Within a level, electrons that penetratemore
  (closer to nucleus) shield theother electrons
  more effectively
• Extent of penetration
• S gt p gt d gt f
• Inner (n1) electrons shield outer (n2)
  electrons very effectively
• 2s electrons spend more time closer to nucleus
  than 2p electrons, thus somewhat shielding the 2p
  electrons

7
Review - Element Properties

• Atomic Size
• Atomic size generally decreases left to right
  across a period of periodic chart
• Atomic size generally increases down a group
• Outer electrons in higher periods (n 1 , 2,
  etc.) lie farther from the nucleus

8
Review - Element Properties

• Ionization Energy (IE)
• Energy required to remove the highest energy
  electron from1 mol of gaseous atoms
• Relative magnitude of IE influences the types of
  bonds
• Element with a low IE is morelikely to lose
  electrons
• Element with a High IE is more likely to gain
  electrons
• IE generally increases left toright (Higher Zeff
  holds electronstighter)
• IE generally decreases down agroup (greater
  distance from nucleus
• Trends in IE are opposite those in atomic size
• Easier to remove low IE electron that is farther
  from nuclues

9
Review - Element Properties

• Electronegativity (EN)
• A number that refers to the relative ability of
  an atom in a covalent bond to attract shared
  electrons
• EN generally increases left to right across a
  period
• Higher Zeff and shorter distance from the nucleus
  strengthen the attraction for the shared pair
• EN generally decreases down a group
• Greater distance from the nucleus weakens the
  attraction for the shared pair
• Trends is Electronegativity are opposite those in
  Atomic size and the same as those in Ionization
  Energy
• The difference in electronegativity (?EN) between
  atoms in a bond greatly influences physical
  chemical behavior

10
Review - Element Properties

• Atomic Size, Ionization Energy, Electronegativity
• Trends in IE are opposite those in atomic size
• Trends is Electronegativity are opposite those in
  Atomic size and the same as those in Ionization
  Energy

11
Review - Element Properties

• Bonds Forces that hold atoms together
• Types of Bonding
• The types of bonding, bond properties, nature of
  orbital overlap, and number of bond determine
  both physical and chemical behavior
• 3 Types Ionic, Covalent, Metallic
• Ionic
• Results from attraction between positive and
  negative ions
• Ions arise through the transfer of electrons
  between atoms with a large ?EN metal
  metalloid non-metal
• Forms crystalline solids with ions packed tightly
  in regular arrays

12
Review - Element Properties

• Types of Bonding
• Covalent
• Results from the attraction between two nuclei
  and a localized electron pair
• Bond arises through electron sharing between
  atoms with small ?EN, usually between 2
  non-metals
• Bond forces include
• Strong covalent bonding forces holding atoms
  together forming a molecule
• Weak intermolecular forces holding separate
  molecules together, thus determining the physical
  properties of covalent compounds
• Produces discrete molecules with specific shapes
  or extended networks of molecules

13
Review - Element Properties

• Types of Bonding
• Metallic
• Results from the attraction between the cores of
  metals atoms (metal cations) and their
  delocalized valence electrons
• The bonding arises through the shared pooling of
  valence electrons from many atoms and leads to
  crystalline solids

Ionic
Covalent
Metallic
14
Review - Element Properties

• Bond Overlap
• Actual bonding in real substances lies between
  the distinct ionic, and decreasingly polar
  covalent models
• Electron Density Maps below showaa
• Small overlap region for the ionic bond in NaCl
• Increase overlap for slightly polar covalent SiCL
  bond in SiCl4
• Highest overlap for non-polar covalent bond in
  ClCl molecule

Ionic
Slightly Polar
Non-Polar
15
Review - Element Properties

• Continuum of Bond Types among Period 3 Elements
• Left Side
• Chlorine compounds display a gradual change from
  ionic to covalent from top to bottom
• Decrease in ionic character (bond polarity) from
  bottom to top
• Right Side
• Elements themselves display a gradual change from
  covalent to metallic from top to botton
• Along the Base
• Compounds of each element display gradual change
  to metallic bonding from left to right
• Decrease in bond polarity (ionic character) from
  left to right

16
Review - Element Properties

• Bond Properties
• There are two (2) important properties of a
  covalent bond
• Bond Length Distance between the nuclei of
  bonded atoms
• Bond Energy (Bond Strength) The Enthalpy change
  (?H) required to break a given bond in 1 mole of
  gaseous molecules
• As bond length increases, bond energy decreases,
  i.e., short bonds are the stronger bonds
• As bond energy decreases, reactivity increases

17
Review - Element Properties

• Orbital Overlap
• In a covalent bond, the shared electrons reside
  in the entire region composed of the overlapping
  orbitals of the two atoms
• Orbitals overlap in two (2) ways
• End-to-End
• s, p, and hybrid orbitals lead to sigma (?) bonds
• Electron density distributed symmetrically along
  bond axis
• Single bond is a ? bond
• Side-to-Side
• p with p (or p with d ) leads to a pi (?) bond
• Electron density distributed above and below bond
  axis
• A double bond consists of one ? bond and one ?
  bond
• A ? bond restricts rotation around the bond axis,
  allowing for different spatial arrangements of
  the atoms, thus, different compounds

18
Review - Element Properties

• Orbital overlap (cont)
• Side-to-Side (cont)
• Pi bonds are often sites of reactivity
• CH2CH2(g) HCl(g) ? CH3-CH2-Cl(g)
• Bond Order ½ the number of electrons shared
• Bond Order 1 single bond
• Bond Order 2 double bond
• Bond Order 3 triple bond
• Fractional bond Order Occurs when a molecule
  has resonance structures for species with
  adjacent single and double bonds

19
Review - Element Properties

• Number of Bonds and Molecular Shape
• The shape of a molecule is defined by the
  positions of the nuclei of the bonded atoms
• The VSEPR theory describes the number of bonding
  and non-bonding electron groups in the valence
  level of a Central atom
• Molecular Notation
• A The Central Atom (Least Electronegative
  atom)
• X The Ligands (Bonding Pairs)
• a The Number of Ligands
• E Non-Bonding Electron Pairs
• b The Number of Non-Bonding Electron Pairs
• Double Triple Bonds count as a single
  electron pair
• The Geometric arrangement is determined by
• sum (a b)

AXaEb
20
Review - Element Properties
AX2E0 a 2 b 0 a b 2 Linear
AX3E0 a 3 b 0 a b 3 Trigonal Planar
AX4E0 or AX2X2E0 a 4 b 0 a b 4 Tetrahedral
AX2E2 a 2 b 2 a b 4 Tetrahedral
AX5E0 a 5 b 0 a b 5 Trigonal Bipyramidal
AX6E0 or AX5X1E0 a 6 b 0 a b 6 Octahedral
AX5N0 or AX4X1E0 a 5 b 0 a b 5 Trigonal Bipyramidal
AX4E1 or AX3X1E1 b 4 a 1 a b 5 Trigonal Bipyramidal
AX6E0 or AX4X2E0 a 6 b 0 a b 6 Octahedral
AX2E3 a 2 b 3 a b 5 Trigonal Bipyramidal
21
Review - Element Properties

• Metallic Behavior
• Elements are often classified as
• Metals
• Metalloids
• Non-metals

22
Review - Element Properties

• Metals, Metalloids, Non-Metals
• Metals lie in the lower-left portion of the
  Period table
• Non-metals lie in the upper-right portion of the
  table
• Metalloids lie between the metals and non-metals
• Intermediate values of
• Atomic size
• Ionization Energy
• Electronegativity
• Shiny solids with low conductivity
• React cation-like (lose e-) with nonmetals
• React anion-like (gain e-) with metals

23
Review - Element Properties

• Metals, Metalloids, Non-Metals
• Metallic Behavior
• Metallic Behavior changes gradually among
  elements
• Metallic behavior parallels atomic size
• Larger members of group (bottom) or period (left)
  are more metallic
• Smaller members are less metallic
• Metals non-metals typically form crystalline
  compounds when they react with each other

24
Review - Element Properties

• Metals, Metalloids, Non-Metals
• Metallic Character
• Ionic size and charge determine the packing in
  the solid
• Ionic size increases down a group
• Ionic size decreases left to right across period
• Cations are smaller than their parent atoms
• Anions are larger than their parent Atoms
• Anions are much large than cations

25
Review - Element Properties

• Acid-Base Behavior
• Oxides are known for almost all elements
• The metallic behavior of an element corresponds
  with the acid-base behavior of its oxide in water
• Acids
• produce Hydrogen (H) ions (Hydronium H3O- ions)
  when dissolved in water
• React with bases to form a salt water
• Bases
• Produces OH- ion in water
• React with acids to form a salt water

26
Review - Element Properties

• Acid-Base Behavior (cont)
• The chart below shows that the electronegativity
  and metallic behavior determine the type of
  bonding between the metal (E) and oxygen (O) in
  the metal oxide
• Elements with low Electronegativity (EN) (metals)
  form basic oxides
• Elements with high EN (non-metals) form acidic
  oxides
• Elements with intermediate EN (some metalloids
  and metals) form amphoteric oxides

27
Review - Element Properties

• Acid-Base Behavior (cont)
• Oxide Acidity
• Increases left to right across a period
• Decreases down a group
• Acidity trends are opposite the trends in
  metallic behavior and atomic size
• When an element forms two oxides, the element has
  a higher oxidation number in the more acidic
  oxide

Increasing Acidity ?
Example SO2 forms weak acid,
H2SO3
(O.N. S 4)whereas SO3 forms strong acid,
H2SO4 (O.N. S
6)
Increasing Acidity ?
28
Review - Element Properties

• Oxidation-Reduction - The relative ability of an
  element to lose or gain electrons when reacting
  with other elements
• Oxidation Number (O.N.) (also called Oxidation
  State)
• O.N. for elements in native state 0
• O.N. the number of electrons that have shifted
  away from the atom (positive O.N.) or toward it
  (negative O.N.)
• An oxidation-reduction (redox) reaction occurs
  when the O.N. values of any atom in the reactants
  are different from those in the products
• All reactions that involve an elemental substance
  (native element) involve an oxidation-reduction
  reaction
• 2K Cl2 ? 2KCl
• All combustion reactions (reactions with elements
  oxygen, i.e., burning)
• CH4 2O2 ? CO2 2H2O

29
Review - Element Properties

• Oxidation is the loss of electrons
• Reduction is the gain of electrons
• Reducing Agent loses electrons (is oxidized,
  attains more positive O.N.)
• Oxidizing agent gains electrons (is reduced,
  attains more negative O.N.)
• Elements with low IE and low EN (groups 1A and
  2A) are strong reducing agent
• Elements with high IE and high EN (/groups 7A and
  oxygen in Group 6A are strong oxidizing agents

30
Review - Element Properties

• Oxidation State of the Main-Group Elements
• Oxidation State A number equal to the magnitude
  of the Charge an atom would have if its shared
  electrons were held completely by the atom that
  attracts them
• The highest (most positive) state in a group
  equals the A group number after all its outer
  (valence) electrons shift toward a more
  electronegative atom
• Among non-metals, the lowest (most negative)
  state equals the A-group number minus 8
• Non-metals have more oxidation states than metals
  in the same group (oxygen and Fluorine are
  exceptions)
• Odd-numbered oxidation states are the most common
  ones in odd-numbered groups
• Even-numbered oxidation states are the most
  common ones in even-numbered groups
• Oxidation states differ by units of 2 because
  electrons are lost or gained in pairs

31
Review - Element Properties

• Oxidation State of the Main-Group Elements
  (cont)
• For many metals and metalloids with more than one
  oxidation state (groups 3A 5A), the lower state
  becomes more common down the group because the np
  electrons only are lost
• An element with more than one oxidation state
  exhibits greater metallic behavior in its lower
  state
• Ex. As3 (lower state) oxide is more basic, more
  like a metal oxide, than is As4 (higher state)

Most common state in Bold
32
Review - Element Properties

• Physical States of Elements
• Physical state solid, liquid, gas and heat of
  phase change vaporization, melting point, etc.
  reflect the relative strengths of the bonding
  and/or intermolecular forces between the atoms,
  ions, or molecules that make up an element or a
  compound
• Metals (lower left of periodic chart)
• Solids
• Strong metallic bonding holds atoms in
  crystalline structures
• Metalloids
• Along staircase line in table and carbon are
  solids
• Strong covalent bonding holds atoms together in
  extensive networks

33
Review - Element Properties

• Physical States of Elements (cont)
• Lighter non-metals and Group 8A (right side and
  top of periodic chart)
• Gases
• Dispersion forces are weak between molecules such
  as H2, N2, O2, F2, Cl2 or atoms with smaller,
  less polarizable electron clouds
• Heavier non-metals
• Liquid (Br2) or soft solids (P4, S8, I2)
• Dispersion forces are stronger between molecules
  with larger, more polarizable electron clouds

34
Review - Element Properties
35
Review - Element Properties

• Phase Changes of the Elements
• Melting Point Boiling Point ?Hfus
  ?Hvap
• Group 1(A)
• These properties generally increase up the group
• The smaller the atomic core, the stronger the
  attraction of the delocalized electrons
• More energy required to melt a solid, boil a
  liquid, etc.
• Groups 7A and 8A
• These properties generally decrease down a group
• Dispersion forces become stronger with the
  larger, more polarizable atoms

36
Review - Element Properties

• Groups 3A to 6A
• These properties reflect changes in interparticle
  forces down the group
• Lower values for molecular non-metals
• Higher values for covalent networks of metalliods
  (and carbon)
• Intermediate values for metals

37
Review - Element Properties

• Physical Properties of Compounds
• Molecular Compounds
• Physical state depends on intermolecular forces
• Polar Compounds
• Dipole-Dipole forces predominate
• Non-Polar Compounds
• Dispersion forces dominate
• Most Molecular compounds are gases, liquids, or
  low melting point solids at room temperature

38
Review - Element Properties

• Physical Properties of Compounds (cont)
• Network Covalent Compounds
• Separate particles absent strong covalent bonds
  link atoms together throughout a network
• Extremely high melting points, boiling points,
  ?Hfus, ?Hvap
• Ex. Silica extended arrays of covalently bond
  silica oxygen atoms
• Ex. Carbon network of covalently bond carbon
  atoms
• Graphite (soft)
• flat sheets of hexagonal carbon rings consisting
  of strong ? bonds and delocalized ? bonds
• Diamond (hardest known substance)
• face-centered cubic cell units

39
Review - Element Properties

• Physical Properties of Compounds (cont)
• Ionic Compounds
• Compounds composed of oppositely charged ions
• Very high Melting Point, Boiling Point, ?Hfus,
  ?Hvap

40
Review - Element Properties

• Physical Properties of Compounds (cont)
• Hydrogen Bonding
• Hydrogen bonding to N, O, F
• Water (H2O) vs Methane (CH4)
• Water
• Hydrogen bonding
• Polar compound
• Much higher MP, BP, ?Hfus, ?Hvap
• Higher specific Heat
• Higher surface tension
• Higher viscosity
• Methane
• No Hydrogen Bonding
• Non-polar compound
• Dominated by dispersion forces