Bonding. Chemical Bond: mutual electrical attraction between nuclei and valence electrons of different atoms

Transcription

1 Chemical Bonding

2 Bonding Chemical Bond: mutual electrical attraction between nuclei and valence electrons of different atoms Type of bond depends on electron configuration and electronegativity

3 Why do atoms bond? TO BECOME STABLE!! Key: 8 valence electrons

4 Types of bonding Ionic: an electrostatic force of attraction between positive and negative ions; ions for when atoms lose/gain electrons Covalent: sharing of electron pairs between atoms polar unequal sharing of electrons nonpolar equal sharing of electrons Metallic: an attraction between metal atoms and outer mobile electrons

5 Ionic Bonding

6 compound formed by the electrostatic force of attraction between positive and negative ions. It involves a transfer of electrons # (+) = # (-) (metals nonmetals) Formula unit: simplest combining ratio of ions in a compound does not exist independently

7 Dot Diagram for NaCl

8 Dot Diagram for CaBr2

9 Warm up Write the compound that forms between the following ions A) Mg and F B) Zinc (positive 2) and nitrogen C) phosphorous and cesium Which of the following are properties of ionic compounds? High melting points shares electrons non electrolyte transfers electron Electrolyte Low melting points

10 Crystal lattice: 3-D crystal structure (arrangement of ions) Lattice energy: energy released when one mole of ionic crystalline compound is formed from gaseous ions

11 Properties high melting point most soluble in water crystalline solid good conductors in liquid or aqueous state hard solids but will fracture

12 Metallic Bonding

13 Metallic bonding: an attraction of metallic atoms for delocalized electrons; electrons roam traveling freely from one atom to another positive ions in a sea of electrons

14 Fun Facts The strength of a metallic bond is determined by number of outer electrons softer metals can be combined with harder metals. (forming alloys) The strongest metals are the transition metals. some d electrons are delocalized, involved in bonding

15 Properties good conductors of heat and electricity shiny (luster) hard, metallic crystals malleable ductile high melting point

16 Covalent bonding

17 Covalent Bonding Occurs between NONMETALS Nonmetals have high electronegativities and want to gain more electrons. They cannot lose valence electrons. In order to bond, nonmetals must then SHARE valence electrons. The atoms share enough electrons to obtain 8 valence electrons (including the shared electrons)

18 Terms to know! Molecular compound: neutral compound consisting of nonmetals covalently bonded in which the electrons are shared Molecule: smallest representative unit of a molecular compound, can exist independently

19 Diatomic Molecules (Memorize these!) Diatomic molecules: molecules consisting of two atoms of the same element. They are ALWAYS found in pairs. Ex: H 2 O 2 F 2 Br 2 I 2 N 2 Cl 2 HOFBrINCl Form a 7 starting at atomic number (Z) 7 and include H

20 Properties low melting & boiling points brittle, dull solids or gases poor conductors of heat and electricity

21 polyatomic ions Polyatomic Ions are covalently bonded group of atoms that have a charge. Examples: Nitrate (NO 3- ), Hydroxide (OH - ), Ammonium (NH 4+ ) Nitrate Structure

22 Let s try it! Calcium and nitrate Sodium and phosphate

23 Definitions Bond length: average distance between nuclei of two bonded atoms (sum of atomic radii) Bond angle: angle between two bonds in a molecule Bond(strength) energy: energy needed to break a bond and form neutral atoms

24 Facts As bond length increases, bond energy decreases. More bonding regions increases the amount of bond energy As electronegativity differences increase, bond energy increases. **Think about it!!!** What relationship do each of the above have?

25 Apply That Information! These diagrams are not drawn to scale Answer: What can you say about the bond strength and bond length of these 2 compounds?

26 Octet Rule Definition: atoms tend to gain, lose or share electrons so that they have eight valence electrons exceptions Hydrogen (2e - ), Beryllium (4e - ), Boron (6e - ) Some elements can have expanded octets (more than 8) example: Sulfur (can have 16 e - )

27 Other covalent bonds Covalent network bonding: 3-D network of covalently bonded atoms macromolecules Very high melting points, very hard solids Ex: diamond, graphite, quartz

28 Network Covalent Bonding

29 Coordinate Covalent bond When both of the shared electrons in a covalent bond come from the same atom

30 Equal sharing of electron between atoms in a compound exist in NON POLAR MOLECULES Unequal sharing of electrons is a POLAR MOLECULE

31 Determine Bond from Electronegativity Differences Non-polar Polar Covalent Ionic

32 Determine the type of bond Use the following electronegativity values to determine the bond character of each: H S (2.2, 2.6) Cs S (0.8, 2.6) S Cl (2.6,3.2) O O (3.4,3.4) Rb Se (0.8,2.6) F I (4.0,2.7) C I (2.6,2.7) Ca N (1.0,3.0)

33 Molecular geometry a.k.a. Molecular Shape

34 Polar vs. Nonpolar Nonpolar molecule: equal sharing of electrons between atoms in a compound, no positive or negative poles exist ex: H 2 H : H Polar molecule: unequal sharing of electrons within a molecule positive and negative ends exist ex: H 2 O

35 Molecular polarity the distribution of molecular charge (even or uneven) Molecular polarity depends on: symmetry molecular shape Lone pairs Molecular polarity influences intermolecular forces.

36 Polar vs Non Polar Structures Water is polar Methane is non polar Hydrogen Fluoride is polar

37 VSEPR theory Valence Shell Electron Pair Repulsion electron pairs (clouds) spread as far apart as possible to minimize repulsive forces

38 Important VSEPR terms Lone Pair: A pair of electrons that do NOT create a chemical bond Shared Pair: A pair of electrons that make a chemical bond between elements Lone Pair of electrons in an electron cloud Shared Pair of electrons

39 VSEPR Shape Linear # Bonds to Central Atom 2 atoms together or 2 bonds to central atom Bent 2 2 # Lone Pairs to Central Atom Examples 0 H-Cl O = C = O 2 1 Bond Angles Trigonal Planar Trigonal Pyramidal Tetrahedral

40 Drawing Structural Formulas 1. Place the least electronegative substance in the middle (Hydrogen will be at the end because it only wants 2 e- s) 2. Calculate the number of valence electrons available (this should always be an even number) 3. Divide that number by 2 to get your bonding pairs 4. Connect all of your atoms in your molecule 5. Use remaining bonding pairs for lone pairs, until all of the bonding pairs are used 6. Check if all atoms are stable(following the octet rule, unless it is an exception), if all happy your structure is done 7. Make double and triple bonds to make remaining atoms stable

41 Let s Try these examples! NH 3 CH 4 H 2 O BF 3 CO 2

42 Molecular Polarity Nonpolar molecules are usually symmetrical. (about their bonds) Polar molecules are usually nonsymmetrical. (about their bonds) Polar molecules are called dipoles. (positive and negative ends) **As soon as you have a lone pair on your central atom, your molecule is polar.** Two shapes are always polar: bent and trigonal pyramidal.

43 Showing molecular polarity To denote a polar bond, an arrow is drawn pointing to the more electronegative substance, with a plus sign at the end that is less electronegative To denote a polar molecule, a lowercase sigma (δ - ) is placed at the more negative end, and a lowercase sigma (δ + ) is placed at the more positive end

44 Intermolecular forces

45 Intermolecular Forces

46 Intermolecular Forces Definitions: weak forces of attraction between molecules Types Dipole-dipole Dipole-induced dipole Hydrogen London Dispersion

47 Dipole-Dipole Exist between polar molecules, higher melting point and boiling point than expected substances exist mostly as solids or liquids

48 Hydrogen bonding Special dipole-dipole force occurring when Hydrogen on one molecule is attracted to N, O, F of another molecule Ex: gives water its unusual properties, ice floats in liquid water, higher melting/boiling point, surface tension

49 Hydrogen Bonding

50 London Dispersion Nonpolar-Nonpolar instantaneous dipole due to a shift in electron strength increase with an increase in number of electrons; low melting point/boiling point, mostly gases Ex: diatomic molecules, noble gases

51 Determine the IMF(s) Present 1) H 2 O 2) SCl 2 3) PF 3 Helpful Hint: Draw the molecular structure and determine if the bond is polar or non polar.

52 Physical properties and bonding Melting point: temperature at which solid liquid Boiling point: temperature at which liquid gas Density: mass/volume, units: g/ml or g/cm 3 Color: some transition metals produce colored ions Ex: copper blue-green Solubility: ability of a solute to dissolve in a given amount of solvent makes a solution

53 How do these properties relate to bond type or intermolecular forces? Stronger bonds/imf s higher melting point/boiling point Weaker bonds/imf s lower melting point/boiling point Solubility likes dissolves like NP-NP no attractive forces exist between solute and solvent, random mixing