Unit 4 Atoms, Bonding, and Chemical Reactions

Ch 18, 19, and 24

CHAPTER 18

Structure of the Atom

• Protons, neutrons, electrons

• Quarks – small particles that make up protons and neutrons

Models

• Dalton - sphere• Thompson – electrons existed• Rutherford – nucleus containing + charge

surrounded by empty space containing electrons• Bohr – electrons travel in orbits around nucleus

with protons and neutrons• Electron Cloud – electrons not in fixed orbits, but

in a cloud around the nucleus• Did The Rabbit Bite Eeyore?

Using the periodic table

• Atomic number = # protons

• Smaller # on periodic table

• On periodic table, # protons = # electrons

• Atomic mass = # protons + # neutrons

• Larger # on period table• # neutrons = mass-

atomic #

• What is the atomic number of zinc?

• How many electrons does tungsten have?

• How many neutrons does scandium have?

• What is the atomic mass of carbon?

• How many protons does astatine have?

Isotopes

• Same number of protons (same element)

• Different number of neutrons

• Therefore, different atomic mass

Periodic Table structure

• Periods (left to right) = increasing number of protons and electrons

• Groups (up and down) = similar reactive properties

Energy Levels

• 1st Level = hold max of 2 e-

• 2nd Level = hold max of 8 e-

• 3rd Level = hold max of 8 e-

Drawing

• Old School Way• Shows all the

electrons on all the energy levels

• Ex: draw flourine

• Electron Dot diagram• Only shows the

outermost electron (the valence electron)

• Ex: draw flourine

Trends

• Left to right, down to up:• Increasing electro

negativity• Increasing ionization

energy• Decreasing atomic radius

Cheats on your periodic table

• On your periodic table, write in the

• Group numbers

• This is the number of electrons in the outermost level

• How many outermost electrons does Boron have?

Bonding

• Atoms want a full outermost shell• This is when they are most stable• Noble Gases (far right of table) already have full

outermost shells• Other elements want to give up or gain e- to

make a full outermost shell• If elements lose an e-, they become positively

charged• If elements gain an e-, they become negatively

charged

Cheats on your periodic tableOxidation numbers

• Group 1 = becomes +1 charged ex: Li+1

• Group 2 = becomes +2 charged ex: Mg+2

• Group 3 = becomes +3 charged ex: Al+3

• Group 5 = becomes -3 charged ex: N-3

• Group 6 = becomes -2 charged ex: O-2

• Group 7 = becomes -1 charged ex: F-1

• Identify the oxidation numbers for each element:

• NaCl

• CaO

• N2O

• SiO2

CHAPTER 19

Ionic vs. Covalent bonding

• Ionic• Total transfer of e-• Between metal and

nonmetal• On both sides of your

stairstep line• Ex: NaCl

• Covalent• Sharing e-• Between a nonmetal

and nonmetal• Both to the right of the

stairstep line• Ex: CO

Identify if it is ionic or covalent

• SiO2

• LiF

• NaCl

• C12H22O11

• HCl

Polar vs. Nonpolar

• Polar• Atoms have diff

electro negativity• Electrons not shared

equally• Ex: HCl• Cl is more

electronegative then H, therefore stronger negative charge

• Nonpolar• Atoms have same

electro negativity• Electrons shared

equally• Ex: Cl2• Same electro

negativity

CHAPTER 24

Chemical Rxns

• Reactants --> Products• Conservation of mass• Mass is converted into different forms but never

created or destroyed

Symbols used in chemical equations

• s solid

• l liquid

• g gas

• aq aqueous, dissolved in water

Coefficients and subscripts

• 4H + 02 2H20

• Notice how this is balanced

• Use the distributive property

• 4 H on left

• 4 H on right

• 2 O on left

• 2 O on right

Chemical Equations

• Balancing equations• Subscripts remain the

same• Coefficient applies to

each element • Ex: 2HO = 2H + 2O

UNL’s tricks to balance!

• 1. Start with compound with the greatest diversity of atoms

• 2. Leave pure elements alone until end (usually O or H)

• 3. If rule #1 doesn’t help, start with the compound farthest left

• 4. All coefficients must be whole numbers. This may require multiplying by the LCM to get rid of fraction.

• 5. # atoms of each element must be balanced on both sides of the equation

Balance these equations

• HgO Hg + O2

• Li + H2O H2 + LiOH

• Mg + O2 MgO

Types of Reactions

• Synthesis A + B --> AB• Ex: 2H2 + O2 2H2O• Decomposition AB --> A + B• Ex: 2H2O 2H2 + O2• Single Displacement A + BC --> AB + C• Ex: Cu + 2AgNO3 Cu(NO3)2 + 2 Ag• Double Displacement AB + CD --> AC + BD• Ba(NO3)2 + K2SO4 BaSO4 + 2KNO3

Energy Exchanges

• Exergonic rxn = releases energy (EXITs)• Ex: glow sticks (releases light)• Exothermic rxn = releases heat• Ex: burning wood• Endergonic rxn = requires energy (moves

IN)• Endothermic rxn = requires heat• Ex: activating a cold pack

Catalysts vs. Inhibitors

• Catalysts• Speed up rxns• Same product is formed• Catalyst remains

unchanged and separate from product

• Enzymes lower the activation E, making the rxn require less E to occur

• Ex: enzymes break down fruit (looks brown)

• Inhibitor• Prevents rxn from

occurring• Same product is formed• Inhibitor remains

unchanged and separate from product

• Ex: lemon juice keeps fruit from browning