unit 4 atoms, bonding, and chemical reactions ch 18, 19, and 24
TRANSCRIPT
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