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Chemical Bonding Ms. Manning.

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Presentation on theme: "Chemical Bonding Ms. Manning."— Presentation transcript:

1 Chemical Bonding Ms. Manning

2 Back to Compounds….. 2 Types: Covalent Compounds
Formed when non-metals bond with other non- metals Ionic Compounds Formed when metals bond with non-metals

3 Conductivity in Liquid
Classification Compound of: Bonding Structure Tm and Tb Conductivity in Solid Conductivity in Liquid Metals Metallic Lattice High Metals and Non Metals Ionic Low Non Metals Covalent Molecular

4 Properties of Metallic compounds
Relatively dense solids (exception Hg) Good conductors of heat and electricity Lustrous when clean/ freshly cut Strong, malleable (can be shaped) and ductile (can be drawn into a wire) Sonorous: Ringing sound when hit Relatively high melting and boiling points Usually form positive ions

5 Properties of Non-Metals
Non-lustrous Can exist in any state - generally gases at room temperature Brittle, non-ductile Poor conductors of heat and electricity Usually exist as molecules in their elemental form Low densities, melting and boiling points. Combine with other nonmetals to form covalent bonds Generally form negative ions, e.g.  Cl-, SO42-, and N3-

6 Properties of Metalloids
Generally look metallic but are brittle (not malleable or ductile) Neither good conductors or insulators; instead they are semiconductors.

7 Chemical Bonding Chemical Bond = the force of attraction holding atoms or ions together This is how compounds are made!

8 Classifying Compounds
Ionic Compound = a pure substance formed from a metal and a nonmetal NaCl CaO Molecular Compound = a pure substance formed from two or more different nonmetals SO2 CO2

9 Ionic versus Molecular Compound
Electrical Conductivity = the ability of a material to allow electricity to flow through it Ionic Compounds  conduct electricity Molecular Compounds  DO NOT

10 Electrolyte Electrolyte = a substance that forms a solution that conducts electricity Ionic compounds  form electrolytic solutions Molecular compounds  form non- electrolytic solutions

11 Ionic Bonding Ions = atoms that have gained or lost electrons
Ionic Bond = the electrostatic attraction between positive and negative ions in a compound Metals lose electrons Non-metals gain electrons Both form octets = MORE STABLE

12 Ionic Bonding – Bohr Diagrams

13 Lewis Dot Diagrams – Ionic Bonding
KBr MgCl2

14 Naming Ionic Bonding Ionic Compounds Metal + Non-metal
Metal name  same as on the atom name Non-Metal  suffix “-ide” Example: NaCl = Sodium Chloride LiF = Lithium Flouride MgO = Magnesium Oxide

15 Non-metal Suffixes Nitrogen = Nitride Oxygen = Oxide
Fluorine = Fluoride Phosphorus = Phosphide Sulfur = Sulfide Chlorine = Chloride Selenium = Selenide Bromine = Bromide Iodine = Iodide Text: page 73 Q

16 How Many Atoms in a Molecule?
Diatomic Molecules = a molecule consisting of two atoms of the same or different elements CO Polyatomic Molecules = a molecule consisting of more than two atoms of the same or different elements NH3

17 Covalent Bonding Covalent Bond = the attractive forces between two atoms that results when electrons are shared by the atoms A simultaneous attraction of two nuclei for a shared pair of electrons In Lewis Diagrams – the shared pairs of electrons are shown as lines and the lone pairs as dots

18 Octet Rule Still Applies!
The shared pair of electrons is considered to be a pair of electrons that make both atoms have an octet 8 8

19 Lewis Dot Diagrams – Covalent Bonds

20 The Lone Pair Lone Pair = a pair of valence electrons not involved in bonding


22 Bonding Capacity Bonding Capacity = the number of electrons lost, gained or shared by an atom when it bonds chemically Allows us to predict how many bonds an atom can form

23 Bonding Capacity Carbon 4 Nitrogen 5 3 Oxygen 6 2 Halogens 7 1
Atom # of Valence Electrons Number of Bonding Electrons Bonding Capacity Carbon 4 Nitrogen 5 3 Oxygen 6 2 Halogens 7 1 Hydrogen

24 Choosing the Central Atom for Polyatomic Molecules
The central position… Is usually occupied by the element with the highest bonding capacity C and N are often in the central position The least electronegative atom is usually the central atom Hydrogen is NEVER the central atom Oxygen and Halogens are usually not the central atom Page 79 gives step by step instructions

25 Covalent Bonds = Strong
A large amount of energy is needed to separate the atoms that make up molecules The stronger the bond the greater the amount of energy needed to break the bond Single bond = strong Double bond = stronger Triple bond = strongest

26 Single, Double & Triple Bonds

27 Polar Covalent Bonds Polar Covalent Bonds = a covalent bond formed between atoms with significantly different electronegativities; a bond with some ionic characteristics When electrons are shared between two atoms = covalent bond In a bond between identical atoms the electrons are shared equally In a bond between two different atoms the sharing is unequal

28 Non-Polar versus Polar Covalent


30 Comparison…

31 Difference in Electronegativity…
If the electronegativity difference is: less than 0.2 = bond is pure covalent is between 0.2 and 1.6 = bond is polar covalent is greater than 1.7 = bond is ionic

32 Polar Molecules Polar Molecules = a molecule that is slightly positively charges at one end and slightly negatively charged at the other because of electronegativity differences

33 Types of Forces Intramolecular Force = the attractive forces between atoms and ions within a compound Ionic Polar Covalent Non-polar Covalent Intermolecular Force = the attractive force between molecules

34 IntRA versus IntER-molecular Forces

35 Some Intermolecular Forces
3 major types of Intermolecular Forces: Dipole-dipole forces London dispersion forces Hydrogen bonding The first two are known as van der Waals Forces London dispersion forces and dipole-dipole forces

36 van der Waals Force Dipole-dipole force = an attractive force acting between polar molecules Attraction between oppositely charged ends of polar molecules

37 van der Waals Force London Dispersion force = an attractive force acting between all molecules including nonpolar molecules A result of temporary displacements of the electron “cloud” around the atoms in a molecule resulting in a extremely short- lived dipole

38 London Dispersion Force

39 Hydrogen Bonding Hydrogen Bonding = a relatively strong dipole-dipole force between a positive hydrogen atom of one molecule and a highly elecgtronegative atom (F, O or N) in another molecule

40 Hydrogen Bonding

41 T M


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