Bonding and Naming.

Bonding and Naming

Bonding and Naming 1. What is a chemical bond?
2. How do atoms bond with each other? How does the type of bonding affect properties of compounds? How can all matter in the universe exist from only 92 elements? Why can you dissolve salt in water but not melt it easily? Why do the carbon atoms in graphite slide off my pencil lead while diamonds are forever? What is the role of carbon in the molecular diversity of life? How can two toxic and violently reactive elements combine to produce table salt (essential for life)? Why is sodium lauryl sulfate found in my shampoo)?

Basics Bond Chemical Bond Why?
something that binds, attaches, or restrains Chemical Bond the force that binds atoms to each other Why? to Achieve Stability (Noble Gas Configuration – s2p6)

Octet Rule Valence electrons Rules
Electrons in the outermost energy level of an atom Rules Octet—Atoms bond to achieve 8 e- in outer energy level Duet—H bonds to achieve 2 e- in outer energy level

Lewis Structures (Electron Dot Diagrams)
Using X to mean any element, draw Lewis Structures for elements in Groups 1,2,13,14,15,16,17,18

x Electron Dot Diagrams Group 17 Group 18 Group 16 Group 14 Group 1

Application of Octet Rule to Ionic Compounds
Note: The element’s symbol represents the core electrons. Dots represent the valence electrons Na · 1s22s22p s1

Application of Octet Rule
Ionic Compounds Metals go to cations Na  Na+ + 1 e- Nonmetals go to anions Cl + 1 e-  Cl- Formation of NaCl [ ] + [ ]- Na Cl

Practice Draw Lewis Structures for the following compounds:
magnesium oxide calcium fluoride aluminum oxide

Answers 2+ [ ]2- O Mg

Answers - - [ ] [ ] 2+ F Ca F

Answers [ ] 2- [ ] O 3+ Al [ ] 2- O [ ] [ ] 3+ Al 2- O

Types of Chemical Bonds
Three basic types of bonds Ionic Electrostatic attraction between ions Covalent and covalent network Shared electrons Metallic Metal atoms bonded to several other metal atoms

Types of Chemical Bonds
Type of bond  properties of compound We will begin with ionic bonding

Ionic Compounds also called salts
held together by attraction between (+) and (-) charges in a lattice (electrostatic attraction) these charged particles are formed by donating (and receiving) electrons have neutral charge overall: the number of + charges = number of – charges strength of the bond is amplified through the crystal structure

Ionic Compounds repeating 3-D structure in crystal lattice = unit cell

Ionic Compounds Energy considerations
Forming a crystal lattice  loss of a lot of PE ↑ stability “lattice energy” = E released when 1 mol of an ionic compound is formed from gaseous ions.

Making NaCl

Lattice Energy Definition:
lattice energy = energy released when a crystal containing 1 mole of an ionic compound is formed from gaseous ions. Related to charge on the ion and distance between nuclei  higher charge  stronger pull  smaller means closer together  stronger pull

Lattice Energies of Some Ionic Compounds
Lattice Energy (kJ/mol) KI -632 KF -808 KBr -671 AgCl -910 RbF -774 NaF NaI -682 LiF -1030 NaBr -732 SrCl2 -2142 NaCl -769 MgO -3795

What happens as the ionic radius increases? (e.g. Check out the Na compounds or the F compounds) What happens as the amount of charge on a single ion increases? Both ions? (look at bottom right hand corner of table)

What happens as the ionic radius increases? Lattice energy decreases – ions are further away from each other. What happens as the amount of charge on a single ion increases? Both ions? Lattice energy increases, especially if both ions have multiple charges.

Predict whether the lattice energy of CsCl is larger or smaller than that of KCl. Predict whether the lattice energy of Na2O will be larger of smaller than that of MgO. What do you think will be the effect of lattice energy on melting point?

Predict whether the lattice energy of CsCl is larger or smaller than that of KCl. Smaller – Cs+ is larger than K+. (CsCl: kJ/mol, KCl: kJ/mol) Predict whether the lattice energy of Na2O will be larger of smaller than that of MgO. Smaller – Na+ is 1+, while Mg2+ has double the charge. (Na2O: kJ/mol) What do you think will be the effect of lattice energy on melting point? The higher the lattice energy, the more energy it takes to melt it, therefore melting point increases as lattice energy increases.

Characteristics solid at room temperature hard, brittle
high melting point, boiling point usually soluble in water conduct electricity when melted or dissolved in water (charges can move) do not conduct electricity when solid

Dissolving NaCl in Water

Naming Chemical Compounds

Terms for describing how atoms bond
Chemical Formula Relative #’s of different elements using symbols and subscripts C2H4 Empirical Formula Chemical Formula Reduced to lowest whole #’s of each element C2H4→CH2 Formula Unit Lowest whole number ratios of Ionic Compounds NaCl, MgBr2, etc.

Terms for describing how atoms bond
Molecular Formula Chemical formula of one molecule e.g. C2H4 Structural Formula of Molecule Shows how atoms are bonded together in a molecule H H C C H H

Naming Ions Monatomic Polyatomic 1 atom + or –
Many atoms covalently bonded with an overall + or – charge (especially anions)

Naming Cations (element’s name + ion) 1. Monatomic Cations
Use Atomic Symbol + Charge e.g. Na+ = sodium Ion (element’s name + ion)

Naming Cations Polyatomic cations (Memorize these!) NH4+ ammonium ion
Hg22+ mercury(I) ion

How Do We Know Charges? Apply Octet Rule (Groups 1, 2, 16, 17)
Memorize Ion List Use Roman numberals to name elements with >1 charge Copper(I) Ion =Cu+ Copper(II) Ion=Cu2+

Can you name the rest of the monatomic anions?
Naming Anions monatomic anion: anions provide the second part of the name of an ionic compound. e.g. Cl- chloride ion use suffix “ide” + “ion” Can you name the rest of the monatomic anions?

Naming Polyatmic Anions
Polyatomic (with oxygens) ite NO2- Nitrite SO32- Sulfite ate NO3- Nitrate SO42- Sulfate ClO-=Hypochlorite=least oxygens ClO2-=Chlorite ClO3-=Chlorate ClO4-=Perchlorate=most Oxygens

Naming Polyatmic Anions
If the anion contains hydrogen, prefix name with “hydrogen” HCO3- = hydrogen carbonate HSO4- = hydrogen sulfate HS- = hydrogen sulfide Let’s try all the phosphates that contain hydrogen:

Naming Binary Ionic Compounds
cation + anion Na Cl- NaCl = sodium chloride

Naming Binary Ionic Compounds
Steps Write Ions Pb4+ O2- Balance Charges (LCM) x2- Empirical Formula/Formula Unit PbO2 Name Lead (IV) Oxide

Practice CaCl2 FeO Fe2S3 Calcium Chloride Iron (II) Oxide
Iron (III) Sulfide

Ionic Compounds Containing Polyatomic Ions
Few positive, many negative polyatomic ions Use Atomic Symbols + Charge polyatomic cation NH4+ ammonium ion Hg22+ mercury(I) ion polyatomic anion OH- hydroxide ion CO32- carbonate ion Note: Charge Refers to whole groups of atoms, not just the last one

Polyatomic Ion Formulae
Write ions NH4+ CO32- Balance Charges (LCM) 2 x Write Formula with lowest whole # ratios (NH4)2CO3 ammonium carbonate Show >1 Polyatomic ion with parentheses

Practice: aluminum sulfate magnesium hydroxide copper(II) acetate

Practice: aluminum sulfate Al2(SO4)3 magnesium hydroxide Mg(OH)2 copper(II) acetate Cu(C2H3O2)2

Oxidation #’s Oxidation #=Charge of an Ion
K+ = oxidation # of +1 O2- = oxidation # of -2 We use oxidation numbers to figure out the formulas of ionic compounds. The sum of oxidation numbers for the formulas of an ionic compounds must = 0.

Metallic Bonds held together by the attraction of free-floating valence electrons (-) for the positive ions in a lattice structure: “electron-sea”

Metallic Bonds What is a regular, repeating three-dimensional arrangement of atoms called? Do the separate electrons that are shown belong exclusively to a single atom? What word is used to describe such electrons? Are the electrons shown the only ones actually present? Explain. Why are the central atoms shown as positively charged? How does the number of separate electrons shown for the group 1A metal atoms compare to the number of atoms? Explain why in terms of valence electrons.

Metallic Bonds - KEY What is a regular, repeating three-dimensional arrangement of atoms called? Crystal lattice Do the separate electrons that are shown belong exclusively to a single atom? What word is used to describe such electrons? no, they are delocalized Are the electrons shown the only ones actually present? Explain. No, they are valence electrons from the metal atoms. Why are the central atoms shown as positively charged? The delocalized (valence) electrons come from the neutral atoms, thus leaving the atoms with a positive charge – i.e. cations. How does the number of separate electrons shown for the group 1A metal atoms compare to the number of atoms? Explain why in terms of valence electrons. They are equal – have only one valence electron.

Metallic Bonds How does the number of separate electrons shown for the group 2A metal atoms compare to the number of atoms? 7. What holds the metal atoms together in such an arrangement? What term is used to describe this model of metallic bonding? How well do metals tend to conduct electricity? How does the model of metallic bonding account for that property? Do metals tend to be brittle, or are they malleable and ductile? How does the model of metallic bonding account for that property?

Metallic Bonds - KEY How does the number of separate electrons shown for the group 2A metal atoms compare to the number of atoms? twice as many electrons What holds the metal atoms together in such an arrangement? Delocalized electrons are simultaneously attracted to > 1 metal cation. What term is used to describe this model of metallic bonding? electron sea How well do metals tend to conduct electricity? How does the model of metallic bonding account for that property? Metals tend to conduct electricity well. The delocalized electrons are not held strongly by individual atoms and are thus able to move easily throughout the metal. Do metals tend to be brittle, or are they malleable and ductile? How does the model of metallic bonding account for that property? malleable and ductile. The delocalized electrons are able to move around the positive metal core atoms and keep the crystal from breaking.

Characteristics of Metals
high m.p. and b.p. Good conductors of electricity and heat conductors in the solid state malleable, ductile shiny, reflective, usually gray (or grey)

Characteristics of Metals
As a metal is struck by a hammer, the atoms slide through the electron sea to new positions while continuing to maintain their connections to each other.

Warmup – Naming Ionic Compounds – binary, monatomic
1. Name the following ionic compounds, and write the pairs of ions that make up these compounds: a) KCl b) SrO c) PbF2

Warmup – Naming Ionic Compounds – binary, monatomic
2. Write the chemical formula for the following ionic compounds, and write the ions that make up these compounds: a) lithium bromide b) barium sulfide c) chromium (III) oxide

Warmup – Naming Ionic Compounds, polyatomic ions
3. Write the chemical formula for the following ionic compounds, and write the ions that make up these compounds: a) cesium oxalate b) calcium hydroxide c) potassium sulfate d) ammonium hydrogen phosphate

Warmup – Naming Ionic Compounds
4. Write the chemical formula for the compounds made up of the following ions, then name the compounds: a) Cs+, O2- b) Pb2+, Br- c) Fe3+, F- d) Na+, CO3-

Ionic Bonding Electron Dot Diagrams
5. How many of each ion will you need to form each of the following ionic compounds? Draw electron dot diagrams to demonstrate the ionic bonding of each compound. a) lithium and bromine b) potassium and sulfur c) aluminum and chlorine d) gallium and oxygen e) bismuth (V) and sulfur

Alloys Mixture of elements that has metallic properties (solid solutions) Substitutional alloys—atoms of similar sizes Brass (Cu +Zn) Bronze (Cu + Sn + Pb) Pewter (Sn + Sb + Pb) Interstitial—much smaller atoms fill spaces between larger atoms Carbon Steel

Alloys It is more difficult for layers of atoms to move over each other, especially in interstitial alloys.

Review Ionic compounds – crystal lattice of oppositely-charged ions, held together by electrostatic charges Metals – lattice of positively-charged ions in a sea of electrons

Covalent Bonding Held together by shared electrons (covalent bond)
Both electrons spend time around each nucleus but spend most of their time in the middle

Diatomic Elements Memorize: N O F Cl Br I H
all exist diatomically in nature (more stable): N2 O2 F2 Cl2 Br2 I2 H2 7- rule: Go to Element #7, travel across to Group 7A, then down in the shape of a 7. Another way to remember them: Hairogens: H2 (H), N2 and O2 (air), F2, Cl2, Br2, I2 (halogens)

Diatomic Molecules 7 Elements bond with themselves (increased stability) H→H2 N→N2 O→O2 F→F2 Cl→Cl2 Br→Br2 I→I2

Covalent Bonding Characteristics of covalently-Bonded compounds (molecules) relatively low m.p. and b.p. do not conduct electricity under any circumstances generally not soluble in water, but soluble in alcohol

Covalent Network Solids
form covalent bonds in all directions  continuous network of strong covalent bonds  no individual molecules extremely hard, very high m.p. and b.p. nonvolatile, insoluble in all solvents brittle, nonconductors of heat and electricity diamond (C), quartz (SiO2)

Compare structures of networks/lattices:
Covalent network (quartz, SiO2) Salt (NaCl) Metal (Cu)

Chemical Bonds Metallic ‘bond’ Ionic bond Covalent bond
Metals and Alloys ‘Sea’ of electrons Ionic bond Ionic Compounds or Salts Metal + Non-metal: NaCl, MgSO4 Electrons exchanged between atoms Covalent bond Molecules and covalent network solids Non-metals: H2O, CH4 Electrons shared among atoms

Formation of a Covalent Bond
When orbitals from two atoms overlap 2 electrons of opposite spin in the overlap As the amount of overlap ↑, the energy of the interaction ↓ When minimum energy is reached, bonding distance occurs Attraction and repulsion of electrons and nuclei are exactly balanced At some distance, nuclei repel, increasing energy again

Formation of a Covalent Bond

Energy Considerations
What does “stable” mean? Changes that lower potential energy are favored. In covalent bonds: Shared electrons  loss of PE  stability Bond energy = energy required to break a chemical bond and form neutral atoms

Relationship between bond length and bond energy in molecules
Bond length = average distance between two bonded atoms (distance of minimum potential energy) As Ebond ↑ , Lengthbond ¯, because the closer the atoms are, the more attraction between nuclei and electron clouds. \ harder to separate.

Naming Molecular Compounds
2 systems – prefixes, vs. oxidation numbers (Honors/AP) See p. 248 in textbook for list of prefixes to memorize

A. Prefixes, roots, suffixes Begin with element with lowest electronegativity. Nitrogen Add appropriate prefix (unless it is mono-). Dinitrogen End second element with "ide" (as for ionic compounds…). Oxide 4. Use appropriate prefix for second element. tetroxide  dinitrogen tetroxide  N2O4

Practice: CCl4 CO CO2 As2S3 P2O5 P4O10

Practice: CCl4 carbon tetrachloride CO carbon monoxide CO2 carbon dioxide As2S3 diarsenic trisulfide P2O5 diphosphorus pentoxide P4O10 tetraphosphorus decoxide

Oxidation Numbers (p. 2 of Honors Supplement)
Oxidation number = the charge on an ion e.g. K+ has an oxidation number of 1+ O2- has an oxidation number of 2- We use oxidation numbers to figure out the formulas of ionic compounds. The sum of oxidation numbers for the formulas of an ionic compounds must = 0.

Oxidation Numbers We can also use oxidation numbers for molecular compounds, by pretending the atoms are ions = “apparent charge”.

Oxidation Numbers The oxidation number of:
an element in the uncombined state is 0. a monatomic ion equals the charge on the ion. hydrogen is generally +1; in hydrides, -1. oxygen is generally -2; in peroxides, -1.

Oxidation Numbers The oxidation number of:
the more electronegative element in a binary covalent compound is negative, while that of the other element is positive. elements other than oxygen and hydrogen in a neutral compound is such that the sum of the oxidation numbers for all atoms in the compound is 0. elements other than oxygen and hydrogen in a polyatomic ion is such that the sum of the oxidation numbers for all atoms in the ion equals the charge on the ion.

Use these rules to assign oxidation numbers to each element in each of the given formulas
e.g. H2O H: 2 x +1, O: -2 N

Assign oxidation numbers to each element:
1. Cl2 2. Cl- Na 4. Na+ 5. KCl 6. H2S CaO 8. H2SO4

Assign oxidation numbers to each element:
10. Cr2O72- 11. NH4Cl 12. NH3 13. NO2 CaH2 (calcium hydride) Na2O2 (sodium peroxide)

Assign oxidation numbers to each element: KEY
1. Cl2 2. Cl- 1- Na 4. Na+ 1+ 5. KCl K: 1+, Cl: 1- 6. H2S H: 2 x 1+, S: 2- CaO Ca: 2+, O: 2- 8. H2SO4 H: 2 x 1+, S: 6+, O: 4 x 2-

Assign oxidation numbers to each element: KEY
10. Cr2O72- Cr: 2 x 6+, O: 7 x 2- 11. NH4Cl N: 5+, H: 4 x 1-, Cl: 1- 12. NH3 N: 3+, H: 3 x 1- 13. NO2 N: 4+, O: 2 x 2- CaH2 (calcium hydride) Ca: 2+, H: 2 x 1- Na2O2 (sodium peroxide) Na: 2 x 1+, O: 2 x 1-

Oxidation Numbers - warmup
Write the oxidation numbers for all elements in the following compounds: P2O5 NO2- NO3- HNO3 K2CrO4 K2Cr2O7

Oxidation Numbers - warmup
Write the oxidation numbers for all elements in the following compounds: P2O5 P: 2 x 5+, O: 5 x 2- NO2- N: 3+, O: 2 x 2- NO3- N: 5+, O: 3 x 2- HNO3 H: 1+, N: 5+, 3 x 2- K2CrO4 K: 2 x 1+, Cr: 6+, O: 4 x 2- K2Cr2O7 K: 2 x 1+, Cr: 6+, O: 7 x 2-

Using Oxidation Numbers to Name Molecular Compounds
Name the following molecular compounds using oxidation numbers: a) CCl4 carbon (IV) chloride b) CO c) CO2 d) P2O5 e) PCl5 f) SO2

Name the following molecular compounds using oxidation numbers: a) CCl4 carbon (IV) chloride b) CO carbon (II) oxide c) CO2 carbon (IV) oxide d) P2O5 phosphorus (V) oxide e) PCl5 phosphorus (V) chloride f) SO2 sulfur (IV) oxide

2. Write formulas for the following molecular compounds: a) carbon(IV) iodide Cl4 b) sulfur(VI) oxide c) nitrogen(IV) oxide d) arsenic(III) sulfide e) phosphorus (III) fluoride

2. Write formulas for the following molecular compounds: a) carbon(IV) iodide Cl4 b) sulfur(VI) oxide SO3 c) nitrogen(IV) oxide NO2 d) arsenic(III) sulfide As2S3 e) phosphorus (III) fluoride PF3

Compounds That Become Acids When Dissolved in Water
General Formula: HX H+ X- monatomic or polyatomic anion

Three Rules: When X ends in “ide” (e.g. chloride, cyanide)  “hydro_______ ic acid” e.g. hydrochloric acid, hydrocyanic acid

Three Rules: When X ends in “ite” (e.g. chlorite, sulfite)  “______ous acid” e.g. chlorous acid sulfurous acid

Three Rules: When X ends in “ate” (e.g. chlorate, sulfate)  “______ ic acid” e.g. chloric acid sulfuric acid

Your turn: HBr HNO2 HNO3

Your turn: HBr hydrobromic acid HNO2 nitrous acid HNO3 nitric acid

Warmup Write formulas for the following molecular compounds:
sulfur trioxide phosphorus pentachloride nitrogen dioxide tetraphosphorus decoxide oxygen difluoride

Warmup Write Lewis structures (dot diagrams) for the following elements: carbon hydrogen fluorine sulfur nitrogen oxygen phosphorus bromine

Molecules and Lewis Structures
Lewis structures show All atoms in the molecule How atoms are connected (# of bonds) Any unshared electron pairs (lone pairs)

Making Lewis Structures
Find total # of valence e-’s for ALL atoms (for ions, consider the charge). 2. Write atom symbols, beginning with the central atom: Carbon in center – CH4 Most electropositive in center – SO42- Non metal (other than H or O) in center – H2PO4- Hydrogen and oxygen are usually on the outside

Making Lewis Structures
Add valence e’-s to each atom, beginning with the bonding electron for each atom added to the central atom. Shared e- pairs represent bonds and are counted in the valences of both elements. 5. Adjust so that each element has 8 valence electrons (2 for H), note exceptions…

Lewis Structure Practice (Electron Dot Diagrams)
Molecules CH3Br BrI H2S PH3 Ions ClO- SO42- H2PO4- NH4+ Hydrocarbons C2H6 C2H4 C2H2 C6H6

Warmup – naming ionic compounds
How many of each ion is in each of the following compounds? AlBr3 PbCl4 RbNO3 MgSO4 K3PO4

How many of each ion is in each of the following compounds? AlBr3 1 Al3+, 3 Br- PbCl4 1 Pb4+, 4 Cl- RbNO3 1 Rb+, 1 NO3- MgSO4 1 Mg2+, 1 SO42- K3PO4 3 K+, 1 PO43-

Write the formula for each of the following compounds: a) ammonium phosphate b) cesium oxide c) copper(I) fluoride d) silver nitride e) beryllium nitrate

Write the formula for each of the following compounds: a) ammonium phosphate (NH4)3PO4 b) cesium oxide Cs2O c) copper(I) fluoride CuF d) silver nitride Ag3N e) beryllium nitrate Be(NO3)2

Exceptions to the Octet Rule
Central atom has < 8 valence e-’s BeF2 BF3 Central atom has > 8 e-’s PF5 SF6 XeF4

Answers C 4 e- Br 7 e- H 1 e- H 1 e- H + 1 e- 14 e- CH3Br

Answers Cl 7 e- O 6 e- + 1 e- 14 e- ClO-

Answers I 7 e- Br + 7 e- 14 e- BrI

Answers 6 e- 6 e- 6 e- 6 e- 6 e- + 2 e- SO42-

Warmup – Lewis Structures
Draw Lewis Structures for the following molecular compounds: carbon tetrachloride sulfite ion ammonia (NH3) water chlorine (Cl2) oxygen (O2) nitrogen (N2)

Molecular Geometry From 2D Lewis Structures to 3D:
Valence Shell Electron Pair Repulsion (VSEPR Theory)

The influence of Unshared Pairs on Geometry
When describing the shape of a molecule, consider the arrangement of the atoms, which is influenced by the unshared (lone) pairs

Molecular Geometry Terminology
Electron domain = region where electron pairs reside = #bonded pairs + #lone pairs Bonding domain = bonded pair Nonbonding domain = lone pair Annotate the diagrams in your notes with these terms

VSEPR Theory – Determining Molecular Shapes
2-3 electron domains

4 electron domains

5 electron domains

5 electron domains (Honors)

6 electron domains

Back to Activity Observe the models, then draw and name the shapes in table B Note: double and triple bonds = 1 lone pair Next, you will build models of the compounds in table A, then draw and name the structures We will discuss polarity later this week.

Hybridization (Honors)
Atomic orbitals are mixed  new, identical hybrid orbitals helps to explain VSEPR # of hybrid orbitals = # atomic orbitals mixed, including lone pairs

Hybridization

Hybrid Orbital Summary
Electron-domain geometry must be known before hybridization is assigned. To assign hybridization: • Draw a Lewis structure. • Assign the electron-domain geometry using VSEPR theory. • Specify the hybridization required to accommodate the electron pairs based on their geometric arrangement. • Name the geometry by the positions of the atoms.

Warmup What is meant by the term “polar”?
Use an example that makes sense to you. Using the PT of theElectronegativities, calculate the difference in electronegativities between a) Na and Cl b) Br and Cl c) B and Cl Identify the bonds in Q2 as ionic, polar covalent or nonpolar covalent.

Warmup What is meant by the term “polar”?
Use an example that makes sense to you. Using the PT of theElectronegativities, calculate the difference in electronegativities between a) Na and Cl (2.1) ionic b) Br and Cl (0.2) NP covalent c) B and Cl (1.0) P covalent Identify the bonds in Q2 as ionic, polar covalent or nonpolar covalent.

Ionic, polar covalent or nonpolar covalent?
The nature of the bond between any two atoms is determined by the difference in their electronegativities (see chart on ho) The greater the difference, the more ionic the bond (on a continuum) ionic covalent 0.4 1.7 4.0 polar (P) NP

Polar covalent range: 1.7 - 0.5 Nonpolar covalent range: 0.4 - 0.0
Ionic range: Cl – Na: = 2.1 (ionic) Polar covalent range: Cl – Al: = 1.5 (polar covalent) Nonpolar covalent range: Cl – Br: = 0.2 (nonpolar covalent) ionic covalent 0.4 1.7 4.0 polar (P) NP Cl - Na Cl - Al Cl - Br

Ionic, polar covalent or nonpolar covalent?
ionic bonds - the less electronegative atom donates 1 or more e-’s to the more electronegative atom covalent bonds - e-’s from both atoms are shared Polar covalent bonds - e-’s are shared, but are not shared equally between two atoms nonpolar covalent bonds- e-’s are shared equally between two atoms

Polar vs. Non-polar: a matter of symmetry

Determining Molecule Polarity
Draw electron dot diagram (Lewis structure) of the molecule, let lines represent bonds Compare the electronegativities of each of the bonded atom pairs Determine whether each bond is polar (P) or nonpolar (NP) Draw an arrow parallel to each bond directed towards the more electronegative atom Is the molecule symmetrical? (If you cut it through the xy, xz and yz planes, would it split into mirror images that are identical?) Yes - molecule is NP No - molecule is P

Determining Molecule Polarity
Examples HCl CCl4 CH3Cl NH3 BF3 H2O

Back to Activity Last column of Data Table B
Last two columns of Data Table A

Sigma Bonds and Pi Bonds (Honors) canceled this year
Sigma (s) bonds: electron density lies on the axis between the nuclei. • All single bonds are s bonds.

Sigma Bonds and Pi Bonds (Honors) canceled this year
• What about overlap in multiple bonds? • Pi (p) bonds: electron density lies above and below the plane of the nuclei. • A double bond consists of one s bond and one p bond. • A triple bond has one s bond and two p bonds (above and below the plane of the nuclei; in front of and behind the plane of the nuclei).

Sigma Bonds and Pi Bonds canceled this year

Formation of a Double Bond canceled this year

Triple Bond canceled this year

Warm Up For each of the following three compounds:
CBr4, CH2O, and CCl2F2 Draw the Lewis structures Name the shape Calculate the polarity of each bond Predict the polarity of each molecule (Honors only) What is the hybridization of the central atom?

Intermolecular forces (van der Waals forces)
Dispersion (London) forces Dipole force Hydrogen bonding

1. Dispersion (London) forces
named after Fritz London, ) a. weakest intermolecular force b. results from the constant motion of electrons  uneven distribution of electrons at any particular moment: “temporary dipole” which may  dipole in nearby molecule. c. acts on all molecules all the time d. only intermolecular force acting among noble gas atoms and nonpolar molecules

1. Dispersion (London) forces
e. with number of electrons: note m.p., b.p. e.g. halogens F2, Cl2 gases at room T Br liquid at room T (more e-’s than F2 and Cl2) I solid at room T (most e-’s)

1. Dispersion (London) Forces

2. Dipole Force (polar molecules)
a. the attraction between two polar molecules: (-) end of one polar molecule attracts the (+) end of another polar molecule b. more polar  stronger dipole force c. closer together  stronger dipole force

2. Dipole Force (polar molecules)

3. Hydrogen bonding a. always involves H
usually involves O, F or N (small, high electronegativity) b. strongest intermolecular force How strong? 5% of the strength of a covalent bond c.  higher b.p. and higher viscosity e.g. H2O

3. Hydrogen bonding

Comparing Bond Types Properties Ionic: salts Metallic
Covalent: molecules Covalent network Description of bond General appearance Malleable vs. brittle Conduct electricity? Under what circumstances? m.p., b.p. State of matter at room temp?

Comparing Bond Types Electrostatic attraction Sea of electrons
Properties Ionic: salts Metallic Covalent: molecules Covalent network Description of bond Electrostatic attraction Sea of electrons Shared electrons Shared electrons in network General appearance crystalline Shiny, prob. gray Range of colors Malleable vs. brittle brittle malleable Prob. brittle Conduct electricity? yes no Under what circumstances? dissolved in H2O or molten all N/A m.p., b.p. High Low Very high State of matter at room temp? Always solid Mainly solid Gas/solid solid

Comparing ALL bond types:
Which is stronger? covalent network > metallic > ionic > covalent (molecules) > H bond > dipole > dispersion e.g. Compare melting points: SiO2 > Fe > NaCl > C12H22O11 > H2O > HCl > H2 sand > iron > salt > sugar > ice > hydrogen chloride > hydrogen gas

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