Unit 3 IB CHEM

Chemical Bonding

• Atoms bond to achieve stability.

• A chemical bond holds chemical compounds together

• Three types of bonding:

  • Ionic: between a metal & a nonmetal or polyatomic

  • Covalent: between two or more nonmetals

  • Metallic: between metals - alloys

Ionic Structure

• Ionic Compound form a large crystal lattice

  • no individual molecules

Covalent Structure

• Covalent compounds form individual molecules

Covalent Compound properties

• Solids are usually soft

  • Ex: Wax

• Typically insoluble in water

• Do NOT conduct electricity     

• Melting points and boiling points are usually low

Covalent Properties (continued)

• Between nonmetals

• Electrons are shared

  • Equal sharing causes non-polar bonds

    • Example:  H₂      

  • Unequal sharing causes polar bonds

Common Substances

• Ammonia - NH₃

• Carbon Monoxide - CO

• Carbon Dioxide - CO₂

• Water -  H₂O

Solid (s) State of Matter

• Fixed volume and shape 

• Particles close together and vibrating in position

• Liquid (l) State of Matter

• Fixed volume, variable shape

  • Takes the shape of its container

• Particles are packed moderately-tightly together

Gas (g) State of Matter

• Variable volume and shape

• Particles far apart and moving very fast

Aqueous (aq) State of Matter

• Fixed volume based on the solution 

• Variable shape

• Particles evenly distributed throughout water

• Ionic compounds dissociate (separate into ions and conduct electricity)

Nonpolar covalent bond

• electrons are shared equally

• Electronegativity difference between 2 bonded atoms = 0.0-0.4

Polar covalent bond

• electrons shared unequally (partial charges)

Electronegativity difference between 2 bonded atoms =0.5-2.1

Ionic Bond

• A bond between a metal and a nonmetal

• Electrons are transferred (full charges)

• Electronegativity difference between 2 bonded atoms = 2.2 - 4.0

Trigonal Planer Bond Angle

120' degrees

Linear Bond Angle

180’ Degrees

Tetrahedral Bond Angle

109.5’ degrees

Trigonal Pyramid Bond Angle

<109.5’ Degrees

Bent VSEPR bond angle

<109.5’ Degrees

Molecular polarity

• For a molecule to be polar it must have a δ+ and a δ- side.

• The molecule must have a non-symmetrical distribution of electrons.

Molecular polarity (how to tell)

• Does the molecule have a lone pair of e- on the central atom?

  • Yes – the molecule is polar

  • No – go to the second question

• Are any of the atoms bonded to the central atom different?

  • Yes – the molecule is polar (different atoms are bonded to the central atom)

  • No – the molecule is nonpolar (all the atoms bonded to the central atom are the same)

Molecular Polarity Arrow

• To indicate direction of polarity draw an arrow with an extra line – the + is on the H (low EN) and the head goes toward the element pulling the e- (N – high EN)

Diatomic Molecules

• Br₂, I₂, N₂, Cl₂, H₂, O₂, F₂

  • Come in pairs

  • More stable this way

Saturated Hydrocarbons

• Saturated

  • no double or triple bonds

• Hydrocarbons 

  • containing hydrogen and carbon

• Are Alkanes

  • generally inert (non-reactive)

  • Due to strength and stability of C-C (348 kJ mol -1) and C-H (412 kJ mol-1) bonds – bond enthalpy 

  • Non polar

Straight chain alkanes

• Single bond between C-C

• 1 Carbon

  • Methane

  • CH₄

Different Representations of Chains

• Molecular Formula 

  • C₃H₈

• Condensed Formula 

  • CH₃CH₂CH₃

• Expanded Structural Formula
  

• Organic Structure

Rules for naming branched alkanes...

Unsaturated Hydrocarbons

• Unsaturated 

  • containing a double or triple bond

• C=C Double bond - Alkenes

• Formula: C_nH_2n

• C≡C Triple Bond - Alkynes

• Formula: C_nH_2n-2

Isomers

• compounds with the same molecular formula but different structural formulas

Naming alkenes with substituents

Properties of Hydrocarbons

• Intermolecular Forces (IMF’s)

  • Nonpolar molecules

    • Only C-H and C-C bonds

• Intermolecular Force

  • London- Dispersion (Van der Waals Forces)

  • Increases with mass 

  • Decreases with branching

Properties of Hydrocarbons (extended)

Homologous Series

• Differ by additional -CH₂-

• Physical Properties:

  • Melting Points/Boiling Points will increase with mass (number of carbons)

• Chemical Properties:

  • Similar

  • Due to similar types of bonds

INTERmolecular

• Attractive forces between a group of molecules

• Weaker than covalent bonds

• Can be broken to change state of matter without changing formula of a compound

INTRAmolecular

• Forces within a molecule

• Strong covalent bonds

• Very strong compared to IMFs

Types of Intermolecular Forces

London Dispersion Force (LDF)

Dipole - Dipole

Hydrodgen Bonding

London Dispersion Force (LDF)

• London dispersion forces (LDF)

  • Occurs between ALL molecules (polar AND nonpolar)

  • Have low melting & boiling points

  • instantaneous temporary dipole

  • larger molecules have greater attractions

  • no partial charges

Dipole Dipole

• Dipole forces (dipole-induced dipole)

  • Stronger than LDF

  • Attraction between permanent polar molecules

  • Partial charges attract each other

Hydrogen Bonding

• Strongest Intermolecular Forces

• Special polar molecules

• Form when Hydrogen is bonded to either:

  • Fluorine

  • Oxygen

  • Nitrogen

• Which makes Hydrogen bonding FON  

Why H-bonds are the strongest

• The strength of the Hydrogen bond is due to the electronegativity difference between hydrogen on one molecule and F, O, or N on a different molecule.

• Remember, F, O, and N are the most electronegative on the periodic table.

Which Molecules have what IMF?

• All Molecules have LDF

• Polar Molecules have both LDF and Dipole - Dipole

• Molecules that have FON bonded with a H have H-bonds as well as both other types

Bonds and forces

• All chemical bonds are strong

  • Atom to atom - Nonpolar and polar covalent compounds (molecules)

  • Ion to ion – ionic compounds

• Ionic Compounds have a network of bonds in a 3-D crystal lattice

• Molecular compounds have weak attractions between the molecules called intermolecular forces

Macromolecular

• Non-metal atoms that are covalently bonded into a lattice structure (e.g., diamond, sand)

• Very rigid structure due to high number of shared electron pairs and covalent bonding throughout structure.  

• No IMFs because there are no small molecules

Molecular polarity

• If the molecule is symmetrical, it is nonpolar

• If the molecule is asymmetrical, it is polar

• Lone pairs of electrons or different elements on the central atom typically make a molecule asymmetrical

IMFs and physical states

• Stronger IMF’s make molecules more attracted to each other

• Stronger attractions take more energy to break and change states of matter

• Remember it can be combination of

  • Dispersion – mass

  • Dipole-dipole – partial charges

  • Hydrogen bonding – very large partial charges

Physical properties of IMFs

• Melting/boiling point, volatility, conductivity and solubility of compounds can be predicted based on IMFs and bond type

• Stronger IMF/attraction = higher MP/BP and lower volatility (evaporates less)

• More free-moving electrons/ions = higher conductivity

• More polar = higher solubility in water, lower solubility in non-polar solvents

• Less polar = higher solubility in non-polar solvents, lower solubility in water

Physical properties of Covalent, Ionic, and metallic

	Covalent Molecules	Ionic Compounds	Metallic
Hardness/ Malleability	Soft and malleable Often gasses	Hard and brittle	Hard and malleable
MP/BP	Low (H2O MP 0ºC,  BP 100ºC)	High (NaCl MP 801ºC BP 1465ºC )	Higher (Fe MP 1538 ºC,  BP 2862 ºC)
Conductivity	None in all states	Good in (l) and (aq) states	Good in (s) and (l) states
Solubility (will or will not dissolve in)	If polar, good in H2O If nonpolar, not good in H2O	Good in water	Only in other metals
Examples	CH3OH         CO2 (polar)     (nonpolar)	NaCl        CaCO3	Fe, Au, Ag, Cu, Al, Pb,...

Not a Flashcard

Physical properties of Covalent, Ionic, and metallic

	Covalent Molecules	Ionic Compounds	Metallic
Hardness/ Malleability	Soft and malleable Often gasses	Hard and brittle	Hard and malleable
MP/BP	Low (H2O MP 0ºC,  BP 100ºC)	High (NaCl MP 801ºC BP 1465ºC )	Higher (Fe MP 1538 ºC,  BP 2862 ºC)
Conductivity	None in all states	Good in (l) and (aq) states	Good in (s) and (l) states
Solubility (will or will not dissolve in)	If polar, good in H2O If nonpolar, not good in H2O	Good in water	Only in other metals
Examples	CH3OH         CO2 (polar)     (nonpolar)	NaCl        CaCO3	Fe, Au, Ag, Cu, Al, Pb,...

Not a Flashcard