Video 3-1 Foundations of Atomic Theory Development of Atomic Models Forces in the Nucleus

Video 3-1

• Foundations of Atomic Theory• Development of Atomic Models• Forces in the Nucleus

Chapter 3

Atoms: The Building Blocks of Matter

I. Foundations of Atomic Theory

Several basic laws were introduced after the 1790’s (emphasis on quantitative analysis):

Law of conservation of mass: mass is neither created nor destroyed during ordinary chemical or physical processes.


Law of definite proportions: chemical compounds contain the same elements in exactly the same proportions by mass regardless of the size of the sample.

Ex. NaCl always is composed of 39.34% sodium and 60.66% chlorine by mass.


Law of multiple proportions: if two or more different compounds are composed of the same 2 elements, the ratio of mass of the second element combined with a certain mass of the first is always a ratio of small whole numbers.


Ex. CO and CO2: For the same mass of carbon, the mass of the O in CO to the mass of O in CO2 will be 1:2

If you had 28 g of CO and 44 g of CO2, both would contain 12 g of C. The CO would contain 16 g of O and the CO2 would contain 32 g of O.


Masses in CO Masses in CO2 Ratios

12 g C16 g O

12 g C32 g O

C= 12:12 = 1:1

O = 16: 32 = 1:2

II. Development of Atomic Models:

John Dalton (1808):1. All matter is composed of extremely

small particles called atoms (cannot be subdivided, created, nor destroyed)

2. Atoms of the same element are identical; atoms of different elements are different


John Dalton (1808):3. Atoms combine in simple whole

number ratios to form compounds4. In chemical reactions, atoms

combine, separate, or are rearranged.


John Dalton (1808):Which of these were later proven wrong

and why?

1) Atoms can be subdivided 2) Atoms of the same element do

NOT have to be identical


J.J. Thomson (1897) and Robert Millikan (1909):

Used cathode rays to determine that atoms contained small negatively charged particles called electrons.

Atoms must also contain positive charges to balance the negative electrons



Other particles must account for most of the mass of the atom

Millikan determine the size of the charge on the electron (oil drop experiment)


Ernest Rutherford (1911): What was the structure of the atom? Gold Foil Experiment

Thomson assumed mass and charged particles were evenly distributed throughout the atom (“plum-pudding” model)



Ernest Rutherford (1911): Expected most of the particles to pass

with only slight deflection Most particles did, but some showed

wide-angle deflections (some almost came back to the source).



Ernest Rutherford (1911): discovery of the NUCLEUS of the

atom small, dense, positively charged center

of the atom number of PROTONS in the nucleus

determines the atom’s identity


Rutherford Atomic Model (solar system model)

III. Forces in the Nucleus

Repulsive forces should exist between protons in the nucleus (like charges repel).

Why doesn’t the nucleus “fly apart” due to the repulsive electromagnetic force?

III. Forces in the Nucleus

Strong (nuclear) force: attractive force that acts over very

small distances in the nucleus causes proton-proton, proton-

neutron, neutron-neutron attractions

Note: gravitational force is present, but negligible. Why?

Video 3-2

• Atomic Dimensions• Properties of Atoms and Ions• Designating Isotopes• Elements on the Periodic Table• Average Atomic Mass

IV. Atomic Dimensions

How “big” are subatomic particles?

Particle Symbol Relative Charge

Mass Number

Relative Mass (amu)

Actual Mass (kg)

electron e- -1 0 0.0005486 9.109 x

10-31

proton p+ +1 1 1.007276 1.673 x

10-27

neutron n0 0 1 1.008665 1.675 x 10-27

e01

p11

n10

IV. Atomic Dimensions

How “big” are subatomic particles? Atomic radii: 40 to 270 pm Nuclear radii: about 0.001 pm Nuclear density: about 2 x 108 metric

tons/cm3

1 amu (atomic mass unit) = 1.660540 x 10-

27 kg Why?

V. Properties of Atoms and Ions atomic number (Z): number of protons in

an atom mass number: number of protons +

neutrons in an atom (number of nucleons—particles in the nucleus)

isotopes (nuclides): atoms of the same element that have different masses (different number of NEUTRONS)

V. Properties of Atoms and Ions ions: atoms with a charge (protons

electrons) charge = protons – electrons atoms can only turn into ions by gaining or

losing ELECTRONS cation: positive ion anion: negative ion

VI. Designating Isotopes There are two ways to write symbols for an

isotope1. name-(mass number)2. symbol

massnumbereratomicnumb

VI. Designating Isotopes

Examples: Hydrogen has 3 isotopes:

Protium Deuterium Tritium

How many neutrons in each isotope? Note: mass number – atomic number =

number of neutrons

Hydrogen-1 H11H21H31

Hydrogen-2

Hydrogen-3

VII. Elements on the Periodic Table

every periodic table will give you at least 3 pieces of information about elements:

Li

3

6.941

Atomic Number

Symbol

Atomic mass (amu)


What is the basis for the atomic mass unit (amu)?

1 amu = exactly 1/12 the mass of a carbon-12 atom (6 protons, 6 neutrons, 6 electrons)

All other atomic masses are based on comparisons to C-12 (exactly 12 amu).


Example:C-13 has a mass that is 1.083613 times

heavier than C-12. The mass of C-13 is

(1.083613) x 12 amu = 13.003356 amu

VIII. Average Atomic Mass Carbon has 3 isotopes (nuclides):

C-12 (12 amu) C-13 (13.003 amu) C-14 (14.003 amu)

Their average mass should be(12 + 13.003 + 14.003) / 3= 13.002 amu

VIII. Average Atomic Mass The atomic mass of carbon (periodic table)

is 12.011 amu. WHY?

VIII. Average Atomic Mass Atomic Mass of an element: weighted

average of all the atoms in a naturally occurring sample of that element (NOTE: not every atom in that sample has the same mass)

Ex. How would you determine the average age of the students in this class?

VIII. Average Atomic Mass atomic mass = sum of (mass of each

isotope x percent abundance)

VIII. Average Atomic Mass Example:

C-12 (12 amu) 98.90% C-13 (13.003 amu) 1.10% C-14 (14.003 amu) trace

atomic mass of C = (12 amu)(0.9890) + (13.003 amu)(0.0110) + (14.003 amu)(0)

= 12.011 amu

VIII. Average Atomic Mass NOTE: the atomic mass of most elements

will usually give you an idea of the most common isotope of that element (mass number that is closest to the atomic mass)

Video 3-3

• Counting Atoms and Stuff

IX. Counting Atoms and Stuff 1 mole = 6.0221367 x 1023 things Ex. 1 mole of eggs contains 6.0221367 x

1023 eggs

6.0221367 x 1023 = Avogadro’s number (N)

1 mol = 6.022 x 1023 particles

IX. Counting Atoms and Stuff 1 mole = 6.0221367 x 1023 things 1 amu = 1.660540 x 10-24 g

Suppose you had a sample of one mole of particles. Each particle weighed exactly 1 amu. How many amu would the sample weigh? How many grams would the sample weigh?

IX. Counting Atoms and Stuff 1 mole of amu = (6.0221367 x 1023)

(1.660540 x 10-24 g)= 1.00000 gram (exactly)

What is the significance of this?

How much will 1 mole of C-12 atoms weigh (in grams)?

12 grams (exactly)

IX. Counting Atoms and Stuff

Molar Mass: mass of one mole of a substance units: grams/mole equal to the ATOMIC MASS of the element for compounds, equal to the SUM of the

masses of all the elements in the compound (multiply each elements’ atomic mass by the subscript)


Example: Find the molar masses of the following:

NaCl CO2

Ca(C2H3O2)2

Na = 22.99 Cl = 35.45 C = 12.01O = 16.00 H = 1.008 Ca = 40.08

58.44 grams/mol

44.01 grams/mol

158.168 grams/mol


1 mole of X = atomic mass of X (grams)

1 mole X = 6.022 x 1023 atoms

These equalities will let you do DIMENSIONAL ANALYSIS Convert grams to moles and moles to grams Convert moles to atoms and atoms to moles

NEVER EVER put 6.022 x 1023 in front of GRAMS


1 molecule C6H12O6 has 6 C atoms, 12 H atoms, and 6 O atoms.

1 mole of C6H12O6 has ___ moles of C atoms, ____ moles of H atoms, and ___ moles of O atoms.

612 6


1 mole XaYb = a moles X = b moles Y

Ex.1 mole Na2C2O4

= 2 moles Na = 2 moles C = 4 moles O

IX. Counting Atoms and Stuff 1 mole X= 6.0221367 x 1023 atoms of X 1 mole of X = atomic mass of X (grams) a moles X = 1 mole XaYb = b moles Y


Moles X Moles XaYb

atoms X

grams Xgrams XaYb

molecules XaYb

1 mole of X = atomic mass of X

1 mole X= 6.0221367 x 1023 atoms

IX. Counting Atoms and Stuff Problem-solving: MAP the problem first (determine what you

are starting with, where you want to end up, and the path to follow).

Example:What is the mass of 3.60 moles of Cu?


Moles X Moles XaYb

atoms X

grams Xgrams XaYb

molecules XaYb


1 mole X= 6.0221367 x 1023 atoms

START

End

IX. Counting Atoms and Stuff Map out the problem first (determine what

you are starting with, where you want to end up, and the path to follow).

Example:What is the mass of 3.60 moles of Cu?

Moles of Cu grams of Cu

3.60 moles Cu xmoles Cu

grams Cu

1

63.546= 229 grams Cu


Map the following problems FIRST, then solve: How many moles are in 11.9 grams? How many atoms are in 3.60 x 10-10 grams

of gold? How many grams of sodium are in 2.34

moles of Na2CO3? How many grams of Fe are in 13.86 grams

of Fe2O3?


Moles X Moles XaYb

atoms X

grams Xgrams XaYb

molecules XaYb


1 mole X= 6.0221367 x 1023 atoms

Documents

Video 3-1 Foundations of Atomic Theory Development of Atomic Models Forces in the Nucleus