thermodynamics study of heat flow

1

Lecture 3

Professor Hicks

Inorganic Chemistry (CHE151)

Coulomb’s Law

• Some types of matter can acquire a

property called charge when rubbed

together

• There are two types

positive (+) negative (-)

+ -

2

Coulomb’s Law

• Opposite charges are attracted

• Like charges are repelled

+ -

+

+ -

-

Discovery of the electron

• Electrons can be made to flow out of all kinds of matter

• They are therefore building blocks of all matter

• They are negatively charged and matter is not charged

there must be a positive part also

• Bending of the rays charge to mass ratio of the electron

3

Millikan oil-drop experiment

Rutherford’s Gold Foil Experiment • particles are positive particles emitted from radioactive substances

• They are also seen flowing towards the negative plate when helium is

placed in the cathode ray tube before evacuating

• Most of the particles go straight through!

The atom must have lots of empty space

Those that are deflected must have hit something so most of the matter of

the atom is compacted into small regions we call the nucleus

4

Subatomic particles

• All matter contains:

negative particles called electrons

positive particles called protons

uncharged particles called neutrons

• Protons and neutrons approximately same

mass (~1.0 amu)

• Electrons mass 1/1836 amu

• Proton and electron have the same exact

charge (1.6 x 10-19 C)

Organization of the atom

• Typical size 10-10 m

• Most of volume

occupied by electrons

• Electrically neutral

#protons = # electrons

electrons

nucleus

10-15 m

10-10 m + charge = - charge

5

Each element has a different

number of protons

# protons

• Same # protons

different # neutron

• Isotopes are versions of an

element with different # neutrons

Isotopes

boron-10

+ +

+

+

+

N

N

N

N

N

boron-11

+ +

+

+

+

N

N

N

N

N

N Both are boron (5 protons)

different # neutrons

different isotopes of boron

= an isotope

6

More on isotopes

Two isotopes of boron

Boron-10

+ +

+

+

+

N

N

N

N

N

Boron-11

+ +

+

+

+

N

N

N

N

N

N

• Elements in nature are mixtures

of different isotopes

• Each Isotope has a different mass

• Nuclei of some isotopes are unstable

- break down in nuclear reactions

Isotope notation

C 13

6 element symbol

nuclear charge (Z)

# protons

mass number (A)

= # protons

+ # neutrons

pronounced “carbon 13”

How many neutrons?

mass number - # protons= # neutrons

= A-Z

=13 – 6 = 7 neutrons

can also be written “carbon-13”

or “C-13”

7

Discuss discovery of the

proton and the neutron

Weighted averages

• A way to calculate different weight for different

types of contribution

Example: Grade is calculated as weighted averages

grade = 15% Quizzes + 60% Exams + 25% Lab

To write it as a math equation convert % /100

grade = 15 x Quizzes + 60 x Exams + 25 x Lab

100 100 100

percentages are the weights

8

Weighted averages

How does it compare to a regular average?

Grade = 15 x Quizzes + 60 x Exams + 25 x Lab

100 100 100

grade = Quizzes + Exams + Lab

3

grade = 33 x Quizzes + 33 x Exams + 33 x Lab

100 100 100

which you can rewrite as

regular average = all same weight

Atomic mass and periodic table

The atomic mass is

weighted average

of isotope masses

# protons

natural Sn is a mixture of:

Sn 120

50

Sn 116

50

Sn 118

50

Sn 122

50

Sn 119

50 Sn 117

50

Sn 115

50 Sn 114

50

Sn 124

50

Sn 115

50 Sn 114

50 Sn 112

50

9

Example. Carbon found in nature is mostly two

isotopes, carbon-12, and carbon-13. They have

abundances of 98.9% and 1.1%. What is the

atomic mass of natural carbon in amu?

atomic mass = 98.9 x 12.00 amu

100

+ 1.1 x 13.00 amu

100

carbon-12 has mass of 12.00 amu

carbon-13 has mass of 13.00 amu

12.01 amu

C 12

6

C 13

6

Mass spectrometer

10

Describe & Discuss

Mass Spectrometer

at the board

Another Example of a Weighted Average

Average Density

• If an object is made up of more than oe type of

matter its density is a weighted average of the

densities of the various components

• If the volumes did not change (as they

sometimes do when homogenous mixtures are

formed)

Average Density = Volume component 1 × Density of component 1

Total volume

+ Volume component 2 × Density of component 2

Total volume

+ …etc.

14

Chlorine consists of two isotopes Cl-35

and Cl-37. Calculate the natural

abundances of each.

15

Ions

• Atoms are electrically neutral

positive charge = negative charge

#protons = # electrons

• Atoms can gain or lose electrons to form

Monatomic ions

• Gaining electrons forms anions (negative ions )

• Losing electrons forms cations (positive ions)

I- K+

lost 2

electrons

Ca2+

N3-

gained 3

electrons

Notating charges

Na+

lost 1

electron K+

lost 1

electron

Al3+

lost 3

electrons

F-

gained 1

electron

gained 2

electrons

O2-

Charge written as number of electrons lost (+) or

gained (-) in superscript

Examples

Na+ K+ Ca2+ Al3+ F- O2- N3-

Ions are visualized

as spheres

16

Trend in type I (fixed-charge) ions

• Same columns (groups) = same charge

• Metals form positive ions

• Non-metals form negative ions

metals

non-metals

metals in this area can have

more than one ion

called variable charge ions

+1

+2 +3 -3

-1 -2

memorize these trends

also a few of these

metals can also form

variable charge ions

Ionic compounds

• Electrically neutral

• Charge positive ions = charge negative ions

NaCl = 1 Na+ and 1 Cl-

Li2S = 2 Li+ and 1 S2-

CaF2 = 1 Ca2+ and 2 F-

-1 +1

-2 +1 +1

2+ -1 -1

Only one compound for

each combination of ions

Na2Cl NaCl2

17

Molecular compounds are more

numerous than ionic compounds • Binary compounds are one with only two

elements

• Knowing what elements are in a binary molecular compound is not enough to determine its chemical formula

N2O4

NO Viagra

activates NO

Rocket fuel

NO2 Smog

N2O

Laughing gas

aka whippets, hippie crack

Formula Unit • Chemical formula of ionic compound called Formula Unit

• Smallest whole number ratio of ions that will be electrically neutral

SrO = 1 Sr2+ and 1 O2-

168 Na+ and 168 Cl-

1019 Na+ and 1019 Cl-

1022 Na+ and 1022 Cl-

CaCl2 = 1 Ca2+ and 2 Cl- Li2O = 2 Li+ and 1 O2-

NaCl = 1 Na+ and 1 Cl-

Na168Cl168

Na Cl 1019 1019

Na Cl 1022 1022

smallest whole number ratio is same for different size lattices

18

Crossing over rule (How to figure out the formula unit of an ionic compound)

Ca Ca2+ N3- N

3-

2+

2+

2+

3-

total +

charge

3 x 2

total -

charge

2 x 3

=

3 2

=

3 2 =

periodic

table

periodic

table

What is the formula unit of the ionic compound

made from calcium and nitrogen?

Why? b/c ionic compounds

electrically neutral

Ca3N2

Crossing over rule

Mg Mg2+ O2- O

2 2

periodic

table

periodic

table


made from magnesium and oxygen?

Why not

Mg2O2?

b/c the formula unit has smallest whole number ratio

of ions that will be electrically neutral (1 to 1 smaller 2 to 2)

MgO

19

Polyatomic ions

• Groups of atoms bonded

together that have a charge

Hg2 2+

PO43- NH4

+ CN-

OH-

SO42-

• Acts as a single ion

H C O O O

atoms

+ e-

H

C O O

O

-1

HCO3-

HCO3-

Polyatomic ions form ionic

compounds

Hg2 2+ PO4

3-

NH4+

CN-

OH-

SO42- HCO3

-

• Positive polyatomic ions can substitute a metal ion

• Negative polyatomic ions can substitute a non-metal ion

20

( ) 2

Crossing over rule (polyatomic ions)

Ca Ca2+ PO43- phosphate

3

periodic

table

table of

ions


made from calcium and phosphate ions?

why use

( ) ? Ca3(PO4)2

(a polyatomic ion)

PO43- PO4

3- Ca2+ Ca2+ Ca2+

2 PO43- ions

not PO423-

Counting atoms in chemical formula

Ca3(PO4)2

PO43- PO4

3- Ca2+ Ca2+ Ca2+

# Ca 2+ ions = 3

# PO43- ions = 2

1 P + 4 O 1 P + 4 O

total = 2 P and 8 O

3 Ca

22

Write formulas for the ionic compounds

formed from:

a) Li+ and O2-

b) Li+ and O22-

c) Hg2+ and F-

d) Hg22+ and F-

e) Hg22+ and O2-

f) Hg22+ and O2

2-

g) Al3+ and C4-

h) Ca2+ and C4-

23

Main Group Elements

Transition Metals

Lanthanides/Actinides

Naming monatomic ions

• Main group negative ions add –ide

Cl- = chloride ion

O2- = oxide ion

• Main group positive ions

“metal name” ion

• Na+ = sodium ion

• Al3+ = aluminum ion

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Naming Type II (variable charge) ions

Fe2+ = Iron (II) ion

Fe3+ = Iron (III) ion

Ag+ = Silver (I) ion

metals

non-metals

most transition metals

can form type II ions

metal name roman numeral of charge ion

example

a few of these

metals also

form type II ions

Naming ionic compounds

(with fixed charge ions)

metal name anion name

sodium bromide

lithium sulfide

NaBr

Li2S

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• metal name anion name

for polyatomic ions

substitute either part

with name Table 8.4

sodium dihydrogen phosphate

lithium dichromate

NaH2PO4

Li2Cr2O7

ammonium sulfate

(NH4)2 SO4

(with polyatomic ions)


• metal name anion name

metal ions that can have different

charges use roman numerals

Iron (III) chloride

Copper (II) dichromate

FeCl3

CuCr2O7

(with variable charge ions)

(Fe3+ and 3Cl-)

(Cu2+ and Cr2O72-)

you will have this

chart on the exam

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Naming ionic compounds with

waters of hydration

• The waters of hydration

are shown after the dot

• Add a prefix for the

number with hydrate

heat

CuSO4 5H2O CuSO4 + 5H2O Copper (II) sulfate pentahydrate

hydrated form anhydrous form

Prefixes

mono, di, tri, tetra,

penta, hexa, hepta,

octa, nona, deca

Naming binary molecular compounds

• Element further to left on PT first

• If in same group lower element comes first

• Both elements get prefixes (mono, di, tri, tetra, penta, hexa, hepta, octa, nona; mono optional)

• Second element gets ending -ide

N2O4

NO nitrogen oxide

dinitrogen tetroxide

NO2 nitrogen dioxide

N2O

dinitrogen oxide

(2 elements)

27

Molecular vs ionic compounds

• Chemical formula of an ionic compound is called its formula unit

• It has smallest whole number ratio of ions

• Formula unit predicted from crossing over rule

making only one formula possible

no prefixes needed (except for waters of hydration)

• In contrast different molecular compounds with same elements are numerous, NO, NO2, N2O, N2O4 etc.

prefixes required to describe each compound

The chemical formulas of molecular compounds

predicted by Lewis Structures ( chapter 10)

thermodynamics study of heat flow

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