C h e m i s t r y 1 A : C h a p t e r 2 P a g e | 1

Chapter 2: Atoms and Elements

Homework:

Read Chapter 2: Work out sample and practice problems in textbook.

Check for the MasteringChemistry.com assignment and complete before due date

Early Ideas on Matter:

Philosophers (Chinese, Greeks, etc) has speculated about the nature of “stuff”

Leucippus (fifth century BC) and his student Democritus (460-370 BC) first

suggested the material world when broken down to the extreme would consist of

indivisible particles called atomos, meaning indivisible.

Alchemists through the middle ages physically experimented with matter aiming to

create gold from base metals and an elixir for everlasting life.

Englishman Robert Boyle (1627-1691) is generally credited as the first to study the

separate science we call chemistry and the first to perform rigorous experiments.

Antoine Lavoisier (1743-1794) discovered the mass of combustion products exactly

equals the mass of the starting reactants.

Law of Mass Conservation (Law of Conservation of Matter); Mass is neither

created nor destroyed in chemical reactions

Joseph Proust (1754-1826) studied copper carbonate, the two tin oxides, and the two

iron sulfides. He made artificial copper carbonate and compared it to natural copper

carbonate, showing that each had the same proportion of weights between the three

elements involved (Cu, C, O). He showed that no intermediate indeterminate

compounds exist between the two tin oxides or the two iron sulfides.

Law of Definite Proportions (Law of Constant Composition); Elements

combine together in specific proportions. All samples of a given compound,

regardless of their source or how they were prepared, have the same

proportions of their constituent elements.

These early ideas led to the foundation steps in atomic theory. Atomic theories explain

the behavior of atoms. We will cover Dalton’s Indivisible atom, JJ Thomson’s Plum

Pudding model, Rutherford’s Nuclear model of the atom, the Bohr’s Quantum (orbit)

model that mathematically only works for one electron systems and the Orbital Wave

Mechanical model. The first three models are found in Chapter 2 while the last two are found in Chapters 7

C h e m i s t r y 1 A : C h a p t e r 2 P a g e | 2

John Dalton (1766-1844) an English chemist that offered evidence to support atoms

using the law of conservation of mass and the law of definite proportions.

Dalton’s Atomic Theory (1808):

1. Elements are composed of tiny, indivisible particles called atoms.

2. Atoms of a given element are identical in properties, but atoms of one

element are different from the atoms of all other elements.

3. Compounds form when atoms of two or more different elements combine in

whole number ratios.

4. Chemical reactions do not create or destroy atoms, they are just rearranged.

Dalton’s atomic theory led to another scientific law…

Law of Multiple Proportions: When two elements form two different

compounds, the masses of elements (B) that combine with 1g of element (A) can be

expressed as a ratio of small whole numbers.

Example: CO(1 g C to 1.33 g O) vs CO2 (1 g C to 2.67 g O)

J. J. Thomson (1856-1940):

Cathode ray tubes (CRT)

contain very low pressures of

a gas and have high voltage

passed through electrodes on

either end. Experiments with

CRT gave off the same

negatively charged radiation

that fluoresced when using

many different gases.

By 1897, JJ Thomson published a paper concluding cathode rays are streams of

negatively charged particles, later known as electrons. This experiment led to a

divisible neutral atom with both negative and positive charges. JJ Thomson called

his atomic theory the Plum Pudding Model of the atom. A positive sphere like

pudding contains particles (plums) of negatively charged electrons that were found

through the CRT experiments. Thomson assumed there were no positively charged

particles since none showed up in the experiment and predicted the mass of the atom

comes from the mass of electrons.

C h e m i s t r y 1 A : C h a p t e r 2 P a g e | 3

In 1909 Robert Millikin succeeded in

measuring the charge of an electron

(1.6022 x 10-19 Coulombs)

through an oil drop experiment

performed numerous times over 5

tedious years. Using Thomson’s

charge to mass ratio (1.7588 x 108

C/g) the electron mass was

accepted as 9.109 x 10-28g, about

2000 times smaller than a single

H atom.

Ernest Rutherford (1871-1937):

In 1910 Ernest Rutherford

created an experiment to test

Thomson’s Plum Pudding model,

the gold foil experiment.

Results showed most heavy

positive alpha particles passed

through a thin gold foil.

Surprisingly, a small portion of

alpha particles were deflected or

even sent back. If Thomson’s

atomic model was correct, this

would be similar to a rifle shot through tissue paper, and no bullet should be

deflected.

The gold foil experiment led to Rutherford’s Nuclear Model of the atom.

The nuclear model has the positive charge (protons) densely set in the center

(nucleus) and the particles of electrons spread out in a cloud around the nucleus.

Neutrons:

It makes no sense to have positive particles (protons) so close together in the

nucleus, they would repel each other. Additionally, mass was missing.

Rutherford’s student, James Chadwick (1891-1974), proposed there are neutrons,

neutral particles within the nucleus similar to protons.

Neutrons were finally isolated in 1932.

C h e m i s t r y 1 A : C h a p t e r 2 P a g e | 4

Atomic Structure:

What we have so far…

Particle Charge Mass (amu) Mass (g)

Electron -1 5.486 x 10-4 amu 9.109 x 10-28 g

Proton +1 1.0073 amu 1.673 x 10-24 g

Neutron 0 1.0087 amu 1.675 x 10-24 g

1 amu = 1.66054 x 10-24g

Atoms are tiny with diameters around 1 - 5 x 10-10 m: 1 angstrom = 1 x 10-10m

Atoms are surrounded by a cloud of negatively charged electrons.

Nucleus contains almost all the mass. It is positively charged and contains

protons (+1 ) and neutrons (0 charge)

The nucleus has a diameter 10,000 times smaller than the atom. This is

equivalent to a marble (nucleus) in the center of a large football stadium.

Neutral atoms have the same number of electrons and protons.

Elements are defined by their number of protons

Electrons can be lost to create a cation or gained to create an anion.

The number of neutrons my vary creating various isotopes.

The unique number of protons is called the Atomic Number (Z)

The protons plus neutrons is the Mass Number (A)

Nuclide symbols: (A Z

X): indicate particular isotopes and ions.

Each element has a unique number of protons (atomic number, Z). The

number of protons defines the element.

Isotopes vary the number of neutrons. Isotopes are chemically identical.

Ions have more or less electrons than protons.

Cations lose electrons, are positive (metals)

Anions gain electrons, are negative (nonmetals)

The nuclide symbol… 𝟏𝟐𝟕 𝟓𝟑

𝐈, represents an iodine atom that has 53 protons,

(127-53) = 74 neutrons and because the charge is neutral it also has 53

electrons. For the iodide ion that has 54 electrons the symbol is…𝟏𝟐𝟕 𝟓𝟑

𝐈-1

C h e m i s t r y 1 A : C h a p t e r 2 P a g e | 5

Example 1:

Fill in the nuclide symbols chart.

Symbol name protons neutrons electrons atomic mass

Barium ion 81 54

5224Cr+3

74 74 184

Carbon-12

Carbon-14

Sulfide ion 16 18

108 47__+1

Finding Patterns: The Periodic Law and Periodic Table:

1869 Dmitri Mendeleev (Russia) and Lothar Meyer (Germany) classified known elements

by organizing similar physical and chemical properties. Mendeleev’s periodic table was an

attempt to organize the known data at the time in a way that made sense. Elements were

arranged by increasing atomic mass and grouped together by chemical reactivity. Several

holes led to predictions of elements and their properties that were not yet discovered “eka-

aluminum” (Ga) and “eka-silicon”(Ge).

Periodic Law – when the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically

Ordered elements by atomic mass

Put elements with similar properties in the same column

Used pattern to predict properties of undiscovered elements

Where atomic mass order did not fit other properties, he re-ordered by other properties Example: Te & I

C h e m i s t r y 1 A : C h a p t e r 2 P a g e | 6

1913 *Henry Moseley improved the periodic table by ordering the elements by increasing

atomic number. More “holes” were found, which led to the discovery of more elements

and the family of noble gases.

The periodic table gives us a great amount of information in an organized manner.

Groups or families are in vertical columns. If properties of a couple elements in a

group are known, one can make a good guess at the properties of other elements

in the same group.

Periods are the horizontal rows in the periodic table.

Many patterns can be seen or predicted following periods and groups. Two more

atomic theories in later chapters will study more of the periodic patterns the theories

attempt to explain.

* Henry Moseley (23 November 1887 – 10 August 1915): In 1910, he started work with

Ernest Rutherford. In 1913, by using x-ray spectra obtained by diffraction in crystals, he

found a systematic relation between wavelength and atomic number, Moseley's law. His

experiments show that cobalt and nickel have clearly differing atomic numbers of 27 and

28, and are correctly placed in the periodic table by an objective measure. Moseley's

discovery showed that atomic numbers have an experimentally measurable basis. In

addition, Moseley discovered gaps in the atomic number sequence at 43, 61, 72, and 75.

These spaces are technetium, promethium, hafnium (discovered 1923) and rhenium

(discovered 1925). In 1914, Henry Moseley enlisted in the Royal Engineers when World

War I broke out. He was killed in action by a sniper in 1915, shot through the head while

telephoning an order. Isaac Asimov once wrote that "In view of what he [Moseley] might

still have accomplished ... his death might well have been the most costly single death of

the war to mankind generally." Many speculated that Moseley should have won the Nobel

Prize, but it is not awarded posthumously. It is speculated that because of Moseley's death

in the War, the British and other world governments began a policy of no longer allowing

their scientists to enlist for combat. Twenty-seven years old at death, Moseley could in

many scientists' opinions have contributed much to the knowledge of atomic structure had

he lived.

C h e m i s t r y 1 A : C h a p t e r 2 P a g e | 7

Atomic Weights: The atomic mass scale has been arbitrarily defined by international agreement. The

scale uses the standard isotope carbon-12, defining its mass to be exactly 12 amu.

Weighted average atomic masses take into consideration the natural abundance of

all the isotopes of an atom.

Masses and isotopic abundances are measured by Mass Spectroscopy.

Difference between simple average and weighted average: Solve for the two averages…

a) Simple average of the numbers… 10.0 g, 16.0 g, 18.0 g, 20.0 g

b) Weighted average given that.

10% of the material is 10.0g, 20% is 16.0g, 30% is 18.0g, 40% is 20.0g

C h e m i s t r y 1 A : C h a p t e r 2 P a g e | 8

Mass Spectrums quantify the result of Mass Spectroscopy…

mass spectrum for zirconium

Number of isotopes

The 5 peaks in the mass spectrum

shows that there are 5 isotopes of

zirconium - with relative isotopic

masses of 90, 91, 92, 94 and 96 on

the 12C scale.

Abundance of isotopes

One can find relative abundances by

measuring the lines on the stick

diagram.

In this case, the 5 isotopes (with their relative percentage abundances) are:

zirconium-90 51.5%

zirconium-91 11.2%

zirconium-92 17.1%

zirconium-94 17.4%

zirconium-96 2.8%

Working out the relative atomic mass

Weighted Atomic Mass = (0.515 x 90)+(0.112 x 91)+(0.171 x 92)+(0.174 x 94)+(0.028 x 96)

= 91.3 (to 3 significant figures) is the relative atomic mass of zirconium.

This example rounds off the isotopic masses much more than generally acceptable

Example 2:

Calculate the weighted average atomic mass of lead given the following information on its

naturally occurring isotopes…

204Pb: 203.973 amu 1.4%

206Pb: 205.974 amu 24.1%

207Pb: 206.976 amu 22.1%

208Pb: 207.977 amu 52.4%

C h e m i s t r y 1 A : C h a p t e r 2 P a g e | 9

Example 3:

There are two naturally occurring isotopes of chlorine. Calculate the percent abundance of

each isotope given the following information on the masses and given that the naturally

occurring weighted atomic mass of chlorine is 35.453 amu… 35Cl: 34.9689 amu

37Cl: 36.9658 amu

Periodic Table Groups:

Metals

Nonmetals

Metalloids/semimetals

Main Group

Transition Metals

Inner Transition Metals or Actinides and Lanthanides

Alkali Metals

Alkaline Earth Metals

Halogens

Noble Gas

Others

Predicting Ion Charges:

Cations: +1, +2, +3, varies

Naming cations:

Anions: -1, -2, -3

Naming anions

C h e m i s t r y 1 A : C h a p t e r 2 P a g e | 10

Counting Atoms by Moles:

Avogadro’s number: 6.022 x 1023 particles = 1 mole

Converting atoms to moles

Converting moles to atoms

Molar Mass:

Solving for molar mass of compounds and molecules

O2

NaHCO3