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Chemical BONDING.

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Presentation on theme: "Chemical BONDING."— Presentation transcript:

1 Chemical BONDING

2 Chemical Bond A bond results from the attraction of nuclei for electrons All atoms trying to achieve a stable octet IN OTHER WORDS the p+ in one nucleus are attracted to the e- of another atom Electronegativity

3 What did the atom of fluorine
say to the atom of sodium? You complete me.

4 Bond Formation exothermic process ENERGY Reactants Energy released

5 Breaking Bonds Endothermic reaction
energy must be put into the bond in order to break it ENERGY Products Energy Absorbed Reactants

6 Bond Strength Strong, STABLE bonds require lots of energy to be formed or broken weak bonds require little E

7 Is a bond forming or breaking? Strong or weak bond?
Reaction Time Energy (KJ) energy absorbed endothermic bond breaking weak unstable bond Products Reactants

8 Is a bond forming or breaking? Strong or weak bond?
Reaction Time Energy Energy absorbed endothermic bond breaking strong stable bond Products Reactants

9 Is a bond forming or breaking? Strong or weak bond?
Energy (KJ) Reaction Time Energy released exothermic bond formation weak unstable bond Reactants Products

10 Is a bond forming or breaking? Strong or weak bond?
Energy (KJ) Reaction Time Reactants Products Energy released exothermic bond formation strong stable bond

11 Two Major Types of Bonding
Ionic Bonding forms ionic compounds transfer of e- Covalent Bonding forms molecules sharing e-

12 One minor type of bonding
Metallic bonding Occurs between like atoms of a metal in the free state Valence e- are mobile (move freely among all metal atoms) Positive ions in a sea of electrons Metallic characteristics High mp temps, ductile, malleable, shiny Hard substances Good conductors of heat and electricity as (s) and (l)

13 It’s the mobile electrons that enable me-tals to conduct electricity!!!!!!

14 IONic Bonding electrons are transferred between valence shells of atoms ionic compounds are made of ions NOT MOLECULES ionic compounds are called Salts or Crystals

IONic bonding Always formed between metals and non-metals [METALS ]+ [NON-METALS ]- Lost e- Gained e-

16 IONic Bonding Electronegativity difference > 2.0
Look up e-neg of the atoms in the bond and subtract NaCl CaCl2 Compounds with polyatomic ions NaNO3


18 Anion (-) Cation (+)

19 Properties of Ionic Compounds
SALTS Crystals hard 22oC high mp temperatures nonconductors of electricity in solid phase good conductors in liquid phase or dissolved in water (aq)

20 Covalent Bonding Pairs of e- are shared between non-metal atoms
molecules Pairs of e- are shared between non-metal atoms electronegativity difference < 2.0 forms polyatomic ions

21 Properties of Molecular Substances
Covalent bonding Low m.p. temp and b.p. temps relatively soft solids as compared to ionic compounds nonconductors of electricity in any phase

22 Covalent, Ionic, metallic bonding?
NO2 sodium hydride Hg H2S sulfate NH4+ Aluminum phosphate KH KCl HF CO Co Also study your characteristics!

23 Drawing ionic compounds using Lewis Dot Structures
Symbol represents the KERNEL of the atom (nucleus and inner e-) dots represent valence e-

24 [Na]+ [ Cl ]- NaCl This is the finished Lewis Dot Structure
How did we get here?

25 Step 1 after checking that it is IONIC
Determine which atom will be the +ion Determine which atom will be the - ion Step 2 Write the symbol for the + ion first. NO DOTS Draw the e- dot diagram for the – ion COMPLETE outer shell Step 3 Enclose both in brackets and show each charge

26 Draw the Lewis Diagrams
LiF MgO CaCl2 K2S

27 Drawing molecules using Lewis Dot Structures
Symbol represents the KERNEL of the atom (nucleus and inner e-) dots represent valence e-

28 Always remember atoms are trying to complete their outer shell!
The number of electrons the atoms needs is the total number of bonds they can make. Ex. … H? O? F? N? Cl? C? one two one three one four

29 Methane CH4 This is the finished Lewis dot structure
How did we get here?

30 Step 1 count total valence e- involved Step 2 connect the central atom (usually the first in the formula) to the others with single bonds Step 3 complete valence shells of outer atoms Step 4 add any extra e- to central atom IF the central atom has 8 valence e- surrounding it . . YOU’RE DONE!

31 Sometimes . . . You only have two atoms, so there is no central atom, but follow the same rules. Check & Share to make sure all the atoms are “happy”. Cl Br H O N HCl

32 O O N N DOUBLE bond TRIPLE bond atoms that share two e- pairs (4 e-)
atoms that share three e- pairs (6 e-) N N

33 Draw Lewis Dot Structures
You may represent valence electrons from different atoms with the following symbols x, , CO2 NH3

34 Draw the Lewis Dot Diagram for polyatomic ions
Count all valence e- needed for covalent bonding Add or subtract other electrons based on the charge REMEMBER! A positive charge means it LOST electrons!!!!!

35 Draw Polyatomics Ammonium Sulfate

36 Types of Covalent Bonds
NON-Polar bonds Electrons shared evenly in the bond E-neg difference is zero Between identical atoms Diatomic molecules

37 Types of Covalent Bonds
Polar bond Electrons unevenly shared E-neg difference greater than zero but less than 2.0 closer to more polar more “ionic character”

38 Place these molecules in order of increasing bond polarity which is least and which is most?
HCl CH4 CO2 NH3 N2 HF a.k.a. “ionic character”

39 non-polar MOLECULES Sometimes the bonds within a molecule are polar and yet the molecule is non-polar because its shape is symmetrical. H C Draw Lewis dot first and see if equal on all sides

40 Polar molecules (a.k.a. Dipoles)
Not equal on all sides Polar bond between 2 atoms makes a polar molecule asymmetrical shape of molecule

41 H Cl + -

42 Space filling model “Electron-Cloud” model
H + -

43 Water is asymmetrical + + H O -

44 Water is a bent molecule

45 W - A - T - E - R as bent as it can be! Water’s polar MOLECULE!
The H is positive The O is not - not - not - not

46 Making sense of the polar non-polar thing
BONDS Non-polar Polar Identical Different MOLECULES Non-polar Polar Symmetrical Asymmetrical

47 IONIC bonds …. Ionic bonds are so polar that the electrons are not shared but transferred between atoms forming ions!!!!!!

48 4 Shapes of molecules

49 Linear (straight line)
Ball and stick model Space filling model

50 Bent Ball and stick model Space filling model

51 Trigonal pyramid Ball and stick model Space filling model

52 Tetrahedral Ball and stick model Space filling model

53 Intermolecular attractions
Attractions between molecules van der Waals forces Weak attractive forces between non-polar molecules Hydrogen “bonding” Strong attraction between special polar molecules

54 van der Waals Non-polar molecules can exist in liquid and solid phases
because van der Waals forces keep the molecules attracted to each other Exist between CO2, CH4, CCl4, CF4, diatomics and monoatomics

55 van der Waals periodicity
increase with molecular mass. Greater van der Waals force? F2 Cl2 Br2 I2 increase with closer distance between molecules Decreases when particles are farther away

56 Hydrogen “Bonding” Strong polar attraction
Like magnets Occurs ONLY between H of one molecule and N, O, F of another H “bond”

57 H is shared between 2 atoms of OXYGEN or 2 atoms of NITROGEN or 2 atoms of FLUORINE Of 2 different molecules

58 Why does H “bonding” occur?
Nitrogen, Oxygen and Fluorine small atoms with strong nuclear charges powerful atoms very high electronegativities

59 Intermolecular forces dictate chemical properties
Strong intermolecular forces cause high b.p., m.p. and slow evaporation (low vapor pressure) of a substance.

60 Which substance has the highest boiling point?
HF NH3 H2O WHY? Fluorine has the highest e-neg, SO HF will experience the strongest H bonding and  needs the most energy to weaken the i.m.f. and boil

61 The Unusual Properties of Water
Unusually high boiling point Compared to other compounds in Group 16

62 Density????

63 H2O(s) is less dense than H2O(l)
The hydrogen bonding in water(l) molecules is random. The molecules are closely packed. The hydrogen bonding in water(s) molecules has a specific open lattice pattern. The molecules are farther apart.

64 The End

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