Let’s have some fun with CHEMISTRYYYY!

Let’s have some fun with CHEMISTRYYYY!

Let’s have some fun with CHEMISTRYYYY!

F. BASIC STRUCTURES OF MATTER

1. ATOMS– basic building blocks of matter– STRUCTURE:• protons (p): positively charged

sub-atomic particle• electron (e-): negatively charged

sub-atomic particle• neutron (n): electrically neutral

F. BASIC STRUCTURES OF MATTER

A ZM CMass Number

Atomic Number

Charge

Atomic Symbol

M = (# of p+) + (# of n0)C = (# of p+) - (# of e-)A = # of p+

Atoms of the same element having the same number of protons (same atomic number) but

different number of neutrons and different atomic mass are called isotopes.

ORBITALS AND QUANTUM NUMBERS

• Quantum numbers: describe certain aspects of the atom

• Orbital: specific distribution of space orientations between electrons

• Electron shell: orbitals with the same value n/1st or principal quantum number

• Subshell: same set of n and l (2nd quantum number)

TYPES OF QUANTUM NUMBERS1. First/Principal (n): energy level of the electron (the

higher, the more energy/size)[for the maximum number of orbitals = n2)

2. Secondary/Azithumal (l = 0 until n-1): sublevel within the energy level = shape of the of the orbital

[s = 0 = spherical, p = 1 = dumbell, d = 2 = clover , f = 3 = more complex]

3. Third/Magnetic (ml): the particular orbital and its space orientation [ s = 1, p = 3, d = 7, f = 10]

4. Fourth/Spin (ms): direction of the spin [+ ½ or – ½ ]

RELATIONSHIP BETWEEN QUANTUM NUMBERS

n l Subshell designation Ml

Number of Orbitals in a subshell (n2)

Total number of electrons

1 0 1s 0 1 2

20 2s 0 1 2

1 2p -1, 0, 1 3 6

3

0 3s 0 1 2

1 3p -1, 0, 1 3 6

2 3d -2, -1, 0, 1, 2 5 10

4

0 4s 0 1 2

1 4p -1, 0, 1 3 6

2 4d -2, -1, 0, 1, 2 5 10

3 4f -3, -2, -1, 0, 1, 2, 3 7 14

PAULI’S EXCLUSION PRINCIPLE• A maximum of two electrons can occupy an orbital (with

different spins)• No two electrons can have the same quantum numbers

ELECTRON CONFIGURATION

PERIODIC LAW

• Physical and chemical properties of elements are the periodic functions of their atomic number

• Elements with similar properties = similar arrangement of outer shell electrons/same group

• Valence electrons: found in the outermost shell of elements in their ground state

PERIODIC TRENDSINCREASES GOING DOWN/DECREASES LEFT TO RIGHT

DECREASES GOING DOWN/INCREASES LEFT TO RIGHT

Atomic/Ionic Radius/SizeElectronegativity

(Ability of an element to attract other electrons to itself in a covalent bond)

Metallic Properties (what makes them metallic – shiny, malleable, etc.)

Ionization Energy(Energy required to remove an

electron from the outermost shell)Electron Affinity

(energy released when an electron is added)

• Molecules: two or more atoms tightly bonded together• Ion: atom with a gained or loss electron (anion: negatively

charged; cation: positively charged)

INTRAMOLECULAR FORCES OF ATTRACTION

1. Covalent bond: between 2 atoms; by electron sharing (nonmetallic + nonmetallic)

2. Ionic bond: between + and – ions; form crystal structures; by electron transfer (nonmetallic + metallic)

3. Metallic bond: between free-floating valence electrons = high conductivity and luster (metallic + metallic)

VAN DER WAAL’S FORCES OF ATTRACTION (IMFA’s)

• Between atoms of different molecules1. London dispersion forces: weakest, between nonpolar

atoms and molecules (nonpolar + nonpolar) (CH3-CH3)2. Dipole-dipole forces: between polar molecules (CO)3. Hydrogen bond: between a hydrogen atom of a polar

molecule + (any) F/O/N atoms in another polar molecule (H20, NH3)

4. Ion-dipole forces: strongest, between polar molecules/ions in solutions (ionic molecules + ionic salts/polar solvents)

WRITING CHEMICAL FORMULAS

• Chemical formulas: denote the number and kinds of atoms in a compound (reacting substances)

1. Represent with symbols (cation first, then anion).2. Indicate oxidation states (to the upper right, put

parentheses for polyatomic ions).3. Write the subscript equal to the oxidation number

of the other element (‘switch’) [omit the ‘1’s’].4. Reduce subscripts to their lowest terms.

FORMULAS TO REMEMBER!

1. FORMULA WEIGHT/MASS or MOLECULAR WEIGHT/MASS

– sum of all atomic weight (AW) (g/mol)

2. PERCENTAGE COMPOSITION

FORMULAS TO REMEMBER!

3. EMPIRICAL AND MOLECULAR FORMULAa.) convert mass in grams to mole ** assume 100 g of any sample so that % composition = value of mass in gb.) get the smallest whole number ratio of the

elements ** divide mole value of each element by the smallest mole value

THIS WILL BE YOUR EMPIRICAL FORMULA (EF)

c.) ** get EFM the same way you get FM

FORMULAS TO REMEMBER!

4. STOICHIOMETRY ** Avogrado’s number = 6.02 x 1023 molecules ** Molar Mass = mass in grams of 1 mole of a substance

CONVERSION

GRAMS MOLES MOLECULES

FORMULAS TO REMEMBER!

4. REACTION STOICHIOMETRY CONVERSION use molar mass of A use molar mass of B

grams of moles of moles of grams ofsubstance A substance A substance B substance B

use coefficients of A and B from the

BALANCED EQUATION

FORMULAS TO REMEMBER!4. SOLUTION STOICHIOMETRY

a.) Molarity (M)

a.) Molality (m)

5. DILUTION

FORMULAS TO REMEMBER!6. GAS LAWS

a.) Boyle’s Law

b.) Charles’ Law

c.) Gay- Lussac’s Law

d.) Combined Gas Law

e.) Avogrado’s Law

f.) Daton’s Law … +

g.) Graham’s Law

h.) Ideal Gas Law

R = 0.0821 L-atm/mol-K

FORMULAS TO REMEMBER!

7. GAS DENSITIES AND MOLAR MASS

At STP,standard temperature: 0°C = 273.15 K = 32°Fstandard pressure: 1 atm = 760 torr = 760 mmHgmolar volume of gas: 1 mol = 22.4 L

This states that ideal gases: 1. the movement of particles is in continuous

random motion 2. pressure of gas in due to the bombardment of

the molecules to the container 3. collisions between or among particles are

elastic 4. the kinetic energy of the system is

proportional to temperature 5. the wide separation between molecules cause

the attractive and repulsive forces to be negligible

KINETIC MOLECULAR THEORY OF GASES

PRACTICE QUESTIONS!

1. A substance composed of two or more elements chemically united is called

a) an isotopeb) an elementc) a mixtured) a compound

2. When energy, like light and heat, is liberated during a reaction such as burning of fuels, this type of reaction occurs.

a) endothermicb) nuclear reactionc) exothermicd) both a and b

3. Chemical action may involve all of the following EXCEPT

a) combination of atoms of all elements to form a moleculeb) breaking down of compounds into elementsc) reacting a compound and an elementd) separation of the components of a mixture

4. When decomposed chemically, 73 g of a sample of HCl produces 71 g of Cl2 and 2 g of H2, while 34 g of H2S sample produces 32 g of S and 2 g of H2. This is an example of:

e) Law of Conservation of Massf) Law of Thermodynamicsg) Law of Multiple Proportionsh) Law of Definite Proportions

5. The theory states that no 2 electrons within the same atom can have the same set of quantum numbers.

a) Hund’s Ruleb) Aufbau’s Principlec) Pauli’s Exclusion Theoryd) Dalton’s Rule

6. Elements with the most stable configuration belong to:

e) transition metalsf) inert gasesg) alkali metalsh) halogen family

7. In the periodic table, which can also indicate the highest energy level or the highest principal quantum number of an atom?

a) groupb) atomic numberc) periodd) atomic mass

8. The total number of orbitals in the fourth energy level is:

e) 4f) 16g) 8h) 18

9. Sodium chloride, most commonly known as table salt, has a pH of:

a) = 7b) < 7c) > 7d) = 0

10. What do you call the solid left after the separation of mixtures?

e) solventf) precipitateg) mother liquorh) supernatant

11. From the given statements below, which is INCORRECT?

a) All atoms of a given element are identical.b) Atoms combine in small whole number ratio.c) All atoms change their chemical identity during a chemical

reaction.d) Elements can combine to form compounds.

12. Isotopes, which are the reasons why elements have no whole mass numbers, have:

e) the same number of protons but different number of electrons

f) the same atomic number but different mass numbersg) the same number of protons but different number of

neutronsh) both b and c

13. The intermolecular forces that exist between the molecules of CCl4, a non-polar compound, is/are:

a) dipole-dipoleb) London forcesc) Hydrogen bondd) both a and b

14. If Magnesium forms a compound with chlorine, what is the formula of the molecule?

e) Mg2Clf) MgCl2

g) Mg2Cl3

h) MgCl3

15. Which of the following statements is FALSE for a neutral atom?

a) number of neutrons is not always equal to the number of protons

b) number of electrons in the atom is equal to the number of protons

c) nucleus has a positive charged) nucleus has a neutral charge

16. The maximum number of electrons that can be placed in the second principal energy level, n=2, of an atom is:

e) 6f) 8g) 4h) 2

17. Why are you advised not to heat a tightly closed vessel?

a) heat increases the volume of the vesselb) pressure might decrease causing the vessel to

implodec) pressure will build up causing the vessel to exploded) heat might be trapped inside the vessel

18. Gases will approach ideality (ideal state) at:e) low pressure and high temperaturef) low pressure and low temperatureg) high pressure and high temperatureh) high pressure and high temperature

19. What should be the last 2 coefficients when the following reaction is balanced?FeS2 + O2 ___ Fe2O3 + ___ SO2

a) 2, 3b) 2, 4c) 2, 6d) 2, 8

20. Gas flows from an area of _____ pressure to an are a _____ pressure.

e) higher-lowerf) equal-equalg) medium-highh) lower-higher

21. What type of reaction occurred in the following:4 NiCO3 + O2 2 Ni2O3 + 4 CO2 + heat?

a) double displacementb) single replacementc) combustiond) decomposition

22. The reaction of ethylene, C2H4(g) and hydrogen chloride, HCl(g) to form ethyl chloride C2H5Cl(g) is an example of what type of reaction?

e) substitutionf) synthesisg) neutralizationh) decomposition

23. A real (or non-ideal) gas will:a) have a higher kinetic energy than ideal gasesb) have a higher boiling than ideal gasesc) not have negligible attractive and repulsive forces

between its moleculesd) conform to the kinetic molecular theory of gases

24. The sum of the partial pressures of gas components is equal to the total pressure of the system. This is stated in which law?

e) Boyle’s Lawf) Combined Gas Lawg) Dalton’s Lawh) Gay-Lussac’s Law

25. Which of the following are not STP conditions for 1 mole of ideal gas?

a) 760 torrb) 273 °Cc) 0°Cd) 22.4 L

26. Which of the following is not stated in the Kinetic Molecular Theory of Gases?

e) temperature of gas is related to the kinetic energy of the molecules

f) pressure is due to the bombardment of molecules to walls of container

g) the larger the molecule of a gas, the greater its kinetic energy

h) collisions are elastic

27. Which has more atoms: 100 mol of Cu, 100 mol of Ag, 100 mol of Au, or 100 mol of Hg?

a) Hgb) Auc) Cud) none of the above

28. How many moles of oxygen atoms are there in 3 moles of Ca3(PO4)2?

e) 8f) 4g) 24h) 36

EXERCISES!ATOM # of n0 # of p+ # of e- C M

72 51

-2 127

n0 = M - A = 127 – 52 = 75

p+ = A =52

e- = p+ – C = 52 – (-2) = 54

EXERCISES!ATOM # of n0 # of p+ # of e- C M

72 51

-2 1274 197

75 52 54

n0 = M - A = 197 – 79 = 118

p+ = A =79

e- = p+ – C = 79 – 4 = 75

EXERCISES!ATOM # of n0 # of p+ # of e- C M

72 51

75 52 54 -2 127118 79 75 4 197

3

M = n0 + A = 72 + 51 = 123

e- = p+ – C = 51 - 3 = 48

EXERCISES!Some AW in amu or g/mol

H = 1.0 Na = 23.0 N = 14.0 C = 12.0 O = 16.0 Al = 27.0

• Find the formula weight of:– H2O = 2 (AWH) + AWO

= 2 (1) + 16 = 18 g/mol– NaOH = AWNa + AWO + AWH

= 23 + 16 + 1 = 40 g/mol– Al(NO3)3

= AWAl + 3 (AWN) + 9 (AWO) = 27 + 3 (14) + 9 (16) = 213 g/mol

EF: CH3

EXERCISES!given: MWcompound = 30 g/mol

80% C ; 20% H (assume 100g)

divide by their AW

• MF = MF = MF = 2

MF: 2 (CH3)

MF: C2H6

ATOMS MASS MOLE RATIOCH

80 g20 g

( 6.67( 20

( 1( 3

by smallest mole value

EXERCISES!Some AW in amu or g/mol

H = 1.0 Na = 23.0 O = 16.0

• 1.5 mol NaOH = ____ g NaOH = ____ molecules NaOH

** FW = 23 + 16 + 1 = 40 g/mol– 1.5 mol NaOH

= 60 g NaOH– 1.5 mol NaOH

= 9 x 1023 molecules NaOH

EXERCISES!Some AW in amu or g/mol

H = 1.0 O = 16.0

** 2 H2 + 1 O2 2 H2O

• 2 mol O = ____ mol HO2 mol O

= 4 mol H2O

• 10 g H = ____ g HO

10 g H = 90 g H2O

** make sure the equation is BALANCED!

from FW from FWfrom balanced equation