Ch02 Final Fix

8/9/2019 Ch02 Final Fix

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John E. McMurry Robert C. Fay

Lecture NotesAlan D. EarhartSoutheast Community College Lincoln, NE

General Chemistry: Atoms First

Chapter 2The Structure and Stability of Atoms

Copyright 2010 Pearson Prentice Hall, Inc.


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Copyright 2010 Pearson Prentice Hall, Inc. Chapter 2/3

Conservation of Mass and the

Law of Definite Proportions

2Hg2HgO

chemical formula

O2+

chemical equation


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Chapter 2/4




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Law of Conservation of Mass: Mass is neither

created nor destroyed in chemical reactions.

HgI2(s)+2KNO3(aq)Hg(NO3)2(aq)+2KI(aq)

4.55 g+2.02g = 6.57 g

Aqueous solutions of mercury(II) nitrate and

potassium iodide will react to form a precipitate ofmercury(II) iodide and aqueous potassium iodide.

3.25 g+3.32g = 6.57 g


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Law of Definite Proportions: Different samples of

a pure chemical compound always contain the

same proportion of elements by mass.

By mass, water is: 88.8 % oxygen

11.2 % hydrogen


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Law of Multiple Proportions and

Daltons Atomic Theory

Law of Multiple Proportions: Elements can combine

in different ways to form different compounds, with

mass ratios that are small whole-number multiples of

each other.

8 grams oxygen per 7 grams nitrogen

16 grams oxygen per 7 grams nitrogen

nitric oxide:

nitrous oxide:


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Daltons Atomic Theory

Law of Multiple Proportions: Elements can combine

in different ways to form different compounds, with

mass ratios that are small whole-number multiples of

each other.


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Daltons Atomic Theory Elements are made up of tiny particles called atoms.

Each element is characterized by the mass of its

atoms. Atoms of the same element have the same

mass, but atoms of different elements have differentmasses.

Chemical combination of elements to make different

chemical compounds occurs when atoms join together

in small whole-number ratios.

Chemical reactions only rearrange the way that atoms

are combined in chemical compounds; the atoms

themselves dont change.


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Atomic Structure: Electrons

Cathode-Ray Tubes: J. J. Thomson (18561940)

proposed that cathode rays must consist of tiny,

negatively charged particles which we now call

electrons.


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Atomic Structure: Electrons

Millikans Oil Drop Experiment


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Atomic Structure: Protons and

Neutrons

Atomic Nucleus: When Ernest Rutherford (1871

1937) directed a beam of alpha () particles at a

thin gold foil, he found that almost all the particles

passed through the foil undeflected. A very small

number, however, (about 1 in every 20,000) were

deflected at an angle and a few actually bounced

back toward the particle source.

Rutherford explained his results by proposing thata metal atom must be almost entirely empty space

and have its mass concentrated in a tiny central

core that he called the nucleus.


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Neutrons

Rutherfords Scattering Experiment


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Neutrons

The mass of the atom is

primarily in the nucleus.


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Atomic Numbers

Atomic Number (Z): Number of protons in an

atoms nucleus. Equivalent to the number of

electrons around the atoms nucleus.

Mass Number (A): The sum of the number of

protons and the number of neutrons in an atoms

nucleus.

Isotope: Atoms with identical atomic numbers butdifferent mass numbers.


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Atomic Numbers

carbon-14

C14

6

atomic number

mass number

carbon-12

C

12

6

atomic number

mass number

6 protons

6 electrons

8 neutrons

6 protons

6 electrons6 neutrons


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Atomic Masses and the Mole

carbon-12: 98.89 % natural abundance 12 amu

carbon-13: 1.11 % natural abundance 13.0034 amu

Why is the atomic mass of the element carbon 12.01 amu?

= 12.01 amu

mass of carbon = (12 amu)(0.9889)+(13.0034 amu)(0.0111)

= 11.87 amu + 0.144 amu


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The Mole


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Mole(mol) - the amount of a substance that

contains the same number of entities as

there are atoms in exactly 12g of carbon-12.

This amount is 6.022x1023. It is called

Avogadros number and is abbreviated as N.

One mole (1 mol) contains 6.022x1023

entities (to four significant figures)


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Molar mass is the mass of 1 mole of in grams

eggsshoes

marbles

atoms

For any element

atomic mass (amu) = molar mass (grams)

1 mole 12C atoms = 6.022 x 1023 atoms = 12.00 g

1 12C atom = 12.00 amu

For any molecule

molecular mass (amu) = molar mass (grams)


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Water

18.02g

CaCO3100.09 g

Oxygen32.00 g

Copper

63.55 g

One mole of common substances.


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Molecular mass (or molecular weight) is the sum of

the atomic masses (in amu) in a molecule.

SO2

1S 32.07 amu

2O +2 x 16.00 amu

SO2 64.07 amu

1 molecule SO2 = 64.07 amu1 mole SO2 = 64.07 g SO2

6.022 x 1023 Molecules of SO2


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Avogadros Number (NA): One mole of any element

contains 6.022 141 x 1023 atoms.

Molar Mass: One mole of any element is the amount

whose mass in grams is numerically equivalent to itsatomic mass.


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1 mole = 28.0855 g

6.022 141 x 1023 molecules = 28.0855 g

Silicon:

Avogadros Number (NA): One mole of any element

contains 6.022 141 x 1023 atoms.

Molar Mass: One mole of any element is the amount

whose mass in grams is numerically equivalent to itsatomic mass.


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Chapter 2/30


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Chapter 2/31

Nuclear Chemistry

Nuclear Chemistry: The study of the properties and

changes of atomic nuclei.

Nuclear Reaction: A reaction that changes an atomic

nucleus.


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Nuclear Chemistry

A nuclear reaction changes an atoms nucleus. A chemical

reaction only involves a change in the way that different

atoms are combined.

Different isotopes of an elements have essentially the

same behavior in chemical reactions, but often have

completely different behavior in nuclear reactions.

The energy change accompanying a nuclear reaction is far

greater than that accompanying a chemical reaction.

Comparisons Between Nuclear and Chemical Reactions


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Nuclear Chemistry

Radioactivity: The spontaneous decay and emission

of radiation from an unstable nucleus.

Radioisotope: A radioactive isotope.


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Chapter 2/34

Nuclear Chemistry


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Nuclear Reactions and

RadioactivityBeta () Radiation

A beta particle is an electron.

Beta particle, -


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Radioactivity

Positron Emission

Gamma () Radiation

A gamma particle is a high-energy photon

A positron has the same mass as an electron but anopposite charge. It can be thought of as a positive

electron.

Positron, +


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RadioactivityElectron Capture

A process in which the nucleus captures an inner-

shell electron, thereby converting a proton to a

neutron.


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Chapter 2/39


Radioactivity


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Nuclear Stability

Every element in the periodic table has at least one

radioactive isotope.

Hydrogen is the only element whose most abundant

stable isotope, hydrogen-1, contains more protons (1)than neutrons (0).

The ratio of neutrons to protons gradually increases,

giving a curved appearance to the band of stability.

All isotopes heavier than bismuth-209 are radioactive,

even though they may decay slowly and be stable

enough to occur naturally.


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Nuclear Stability

These processes increase the neutron/proton ratio:

NeutronElectron capture: Proton + Electron

This process decreases the neutron/proton ratio:

Proton + -Beta emission: Neutron

Proton + +

Alpha emission:

NeutronPositron emission:

XZ

A

YZ - 2

A - 4

+ He2

4

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Ch02 Final Fix