chapter 4: atoms and elements - moorpark college

C h e m i s t r y 1 2 C h 4 : A t o m s a n d E l e m e n t s P a g e | 1

Chapter 4: Atoms and Elements

Remember to keep up with MasteringChemistry, Workshops, Mini-Reports and Labs

Early Ideas on Matter:

Philosophers (Chinese- yin/yang; Greeks-earth/wind/fire/water) speculated about

the nature of “stuff” without relying on scientific evidence

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

tiny 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, J.J. 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 4 while the last two are found in Chapters 9.

http://en.wikipedia.org/wiki/Copper

http://en.wikipedia.org/wiki/Cu

http://en.wikipedia.org/wiki/C

http://en.wikipedia.org/wiki/O


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.

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 element (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);

By the mid-1800’s new experiments

gave data that was inconsistent with an

indivisible atom. 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 radiation

that is negatively charged. The same

negative charged substance that

fluoresced (gave off light) was found

using many different gases.

By 1897, JJ Thomson published a paper that concluded the cathode rays are streams

of negatively charged particles, later known as electrons. These particles are smaller

than the atom itself, therefore these are the first sub-atomic particles identified

through experiment.


Plum Pudding Model;

This experiment led to a divisible neutral atom

which must have 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.

Since the atom is neutral, there must be a

positively charged electric field as well.

Thomson assumed there were no positively

charged particles since none showed up in the

experiment. He incorrectly predicted much of the mass of the atom comes from the

mass of electrons.

In 1909 Robert Millikin;

Robert Millikin measured 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.

This caused the

question: What is

the major

contributor of an

atom’s mass.


Ernest Rutherford (1871-1937):

In 1910 Ernest Rutherford created the gold foil experiment experiment to test

Thomson’s Plum Pudding model. The

results showed most of the heavy

positive alpha particles passed right

through a thin gold foil. Surprisingly,

a small portion of alpha particles were

deflected or even sent back. If

Thomson’s Plum Pudding atomic

model was correct, this would be

similar to a rifle shot through tissue

paper, and no bullet should be

deflected.

Rutherford’s Nuclear Model of the

atom explains why some of the alpha particles

were deflected of bounced back as the picture

shows. The nuclear model has all the positive

charge (protons) densely set in the center

(nucleus) and the particles of electrons spread out

in a cloud around the nucleus.

1. Most of the atom’s mass and all of its

positive charge are contained in a small

core called the nucleus.

2. Most of the volume of the atom is empty space through which the tiny,

negatively charged electrons are dispersed.

3. The number of negatively charged electrons outside the nucleus is equal to

the number of positively charged particles (protons) inside the nucleus, so

that the atom is electrically neutral.

The dense nucleus makes up more than 99.9% of the mass of the atom, but

occupies only a small fraction of its volume. The low mass electrons are

distributed through a much larger region. A single grain of sand composed of

just atomic nuclei would have a mass of 5 million kg. Astronomers believe that

black holes and neutron stars are composed of this kind of incredibly dense

matter.


Neutrons: It did not make sense to have all the positive particles (protons) so close

together in a nucleus, they would repel each other. Additionally, some

mass was missing. One of Rutherford’s students, James Chadwick (1891-

1974), proposed there are neutrons, neutral particles within the nucleus

similar to protons. Neutrons were isolated later in 1932.

Atomic Structure:

What we have so far…

Particle Charge Mass (amu) Mass (g)

Electron -1 0.0005486 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

Solve for the inverse of this number: amu = 1 g

Atoms are extremely tiny with diameters around 10-10 m:

The tiny atomic nucleus is surrounded by a very large cloud of electrons

The nucleus contains almost all the mass of an atom. It is positively charged

and contains protons (+1 ) and neutrons (0 charge)

Size Example: a marble (nucleus) in the center surrounded by a large

football stadium (a cloud of electrons).

Neutral atoms have the same number of electrons and protons.

The number of protons defines the element. Each chemical element (X) has a

unique number of protons (atomic number, Z).

Ions have more or less electrons than protons.

Cations lose electrons, are positive (metals)

Anions gain electrons, are negative (nonmetals)

Isotopes will be the same element with the same number of protons, but the

number of neutrons are different. Isotopes are chemically identical.

The protons plus neutrons is the Mass Number (A)

Nuclide symbols (or Isotope Symbol Notation) indicate particular isotopes

and ions. A

Z X


Example 1: Isotope Symbols

Fill in the nuclide symbols chart.

Nuclide symbol, A

Z X

name protons neutrons electrons atomic mass

𝟏𝟐𝟔

C carbon-12 6 6 6 12

carbon-14 6

𝟏𝟒𝟕

N

Sulfide ion 16 18

Potassium ion 18 39

Elements:

Many of our element symbols are based on its English name

Some element names are based on Greek or Latin origins

sodium, Na is from the Latin word natrium

lead, Pb is from the Latin word plumbum which means heavy

phosphorus, P is from the Greek words phôs (light) and phoros

(bearer), Phosphoros was a god of light in Greek myth.

Some element names honor locations or people

Berkelium, Bk, for the location of the lab that created it first.

Einsteinium, Es, for Albert Einstein

Periodic Table: Patterns and the Periodic Law

Development:

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

known elements (about 65 known at that time) and noted similar physical and

chemical properties were found periodically when arranged by increasing

atomic weight 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


The first periodic table:

Ordered elements by atomic mass Put elements with similar properties in the same columns Used pattern to predict properties of undiscovered elements Where atomic mass order did not fit other properties, Mendeleev re-

ordered by other properties Example: Te & I

1915 Henry Moseley developed the concept of atomic numbers. He

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.

Vertical columns are called groups or families. If you are aware of the properties of a

couple elements in a group, you can make a good guess at the properties of the 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. It is easy to identify

certain expected characteristics by locations, such as Metals, Nonmetals Metalloids.

Periodic Table:

Organization:

family/group

period

metals

nonmetals

metalloids/semiconductors,

Groups:

Main Group

Transition Metals

Inner Transition Metals or Actinides and Lanthanides

Alkali Metals

Alkaline Earth Metals

Halogens

Noble Gas

Coinage metals


Atoms may lose or gain electrons to form Ions: Cations (positive charge): Na+

Anions (negative charge): Cl-

Naming: Preview of Chapter 5: Cations with known oxidation state of metal

Group 1A (+1), 2A (+2) , Al and Ga (+3), Zn and Cd (+2), Ag (+1)

Name of ion is identical to the name of the atom for cations

Variable oxidation state of metal

Transition metals and metals below the nonmetal on the right have a

variable oxidation state that must be indicated by Roman Numerals in

parenthesis (this method is what I expect you to learn.

Fe+3, iron (III); Fe+2, iron (II); Cu+1, copper (I); Sn+4, tin (IV)

An alternative method differentiates from the higher oxidation number

and lower oxidation number using the old form of the name and ic or

ous as an ending respectively. (you are not responsible for knowing the

ic and ous ending of metal cations) Fe+3, ferric Fe+2, ferrous; Cu+2,

cupric; Cu+1, cuprous; Sn+4, stannic; Sn+2, stannous

Anions

Naming:

Group VA (-3); VIA (-2), VIIA (-1)

Name of the element root followed by ide.

N-3, nitride; S-2, sulfide, Br-1, bromide

Atomic Weights:

The atomic mass scale is arbitrarily defined by international agreement and is

based a 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.


Mass Spectrum quantifies the results…

The mass spectrum for

zirconium

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.

The abundance of the isotopes

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

(This simple example rounds off the mass much more than I generally accept.)


Working out the relative atomic mass

Using the equation

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 is the relative atomic mass of zirconium.

Mass: Simple vs. Weighted Average:

Simple average: add all the numbers and divide by the count

Solve for the simple average...

Given: 12.0 g, 16.0 g, 17.0 g

Weighted average: Mass = ∑ individual mass value x fractional abundance

Weighted average takes into consideration the fractional abundance of each

number. Fractional abundance is the decimal form of the percent abundance.

All fractional abundance values add up to a total of one (1.00) so there is no

reason to divide by the count.

Solve for the weighted average…

Given: 12.0 g (80.0%), 16.0 g (15.0%), 17.0 g (5.0 %)

Naturally occurring weighted masses for elements are found on the periodic table:

Atomic mass = isotopic mass x fractional abundance

Atomic mass = (massA x fract. abund.A) + (massB x fract. abund.B) + (….

All the fractional abundance values add up to a total of one (1.00)


Example 2:

Use the weighted average to solve the average atomic mass found in nature for Si

Given the following information on its naturally occurring isotopes…

Keep appropriate significant figures.

28Si: 27.977 amu 92.21%

29Si: 28.976 amu 4.70%

30Si: 29.974 amu 3.09%

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 (1-x)

37Cl: 36.9658 amu (x)

Preview: 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 molecules and compounds

O2

H2O

CoBr3

chapter 4: atoms and elements - moorpark college

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