Chapter 3

Atoms: The Building Blocks of Matter

When you crush a lump of sugar, you can see that it is made up of many smaller pieces of sugar. You may grind

these particles into a very fine powder, but each tiny piece is still sugar. Now suppose you dissolve the sugar

in water. The tiny particles seem to disappear completely. Even if you look at the sugar-water solution

through a powerful microscope you cannot see any sugar particles. Yet if you were to taste the solution, you’d know that the sugar is still there. Observations like

these led early philosophers to ponder the fundamental nature of matter. Is it continuous and infinitely divisible, or is it divisible only until a basic, invisible particle that

cannot be divided further is reached?

The Atom: From Philosophical Idea to Scientific Theory

The particle theory of matter was supported as early as 400 b.c. by certain Greek thinkers, such as Democritus. He called nature’s basic particle

an atom, based on the Greek word meaning “indivisible.” Aristotle was part of the generation

that succeeded Democritus. His ideas had a lasting impact on Western civilization, and he did not believe in atoms. He thought that all matter was continuous, and his opinion was accepted

for nearly 2000 years. Neither the view of Aristotle nor that of Democritus was supported by experimental evidence, so each remained

speculation until the eighteenth century. Then scientists began to gather evidence favoring the

atomic theory of matter.

Foundations of Atomic Theory

Nearly all chemists in late 1700s accepted the definition of an element as a substance that cannot be broken down further

Knew about chemical reactions

Great disagreement as to whether elements always combine in the same ratio when forming a specific compound

Law of Conservation of Mass

With the help of improved balances, investigators could accurately measure the masses of the elements and compounds they were studying

This lead to discovery of several basic laws

Law of conservation of mass states that mass is neither destroyed nor created during ordinary chemical reactions or physical changes

Law of Definite Proportions law of definite

proportions A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound

Each of the salt crystals shown here contains exactly39.34% sodium and 60.66%

chlorine by mass.

Law of Multiple Proportions Two elements sometimes combine to form more than one

compound For example, the elements carbon and oxygen form two

compounds, carbon dioxide and carbon monoxide

Consider samples of each of these compounds, each containing 1.0 g of carbon

In carbon dioxide, 2.66 g of oxygen combine with 1.0 g of carbon

In carbon monoxide, 1.33 g of oxygen combine with 1.0 g of carbon

The ratio of the masses of oxygen in these two compounds is exactly 2.66 to 1.33, or 2 to 1

1808 John Dalton Proposed an explanation for the law of conservation of

mass, the law of definite proportions, and the law of multiple proportions

He reasoned that elements were composed of atoms and that only whole numbers of atoms can combine to form compounds

His theory can be summed up by the following statements

Dalton’s Atomic Theory1. All matter is composed of extremely small particles called

atoms.

2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.

3. Atoms cannot be divided, created or destroyed.

4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds.

5. In chemical reactions, atoms are combined, separated, or rearranged.

Modern Atomic Theory Dalton turned Democritus’s idea into a scientific theory

which was testable

Not all parts of his theory have been proven correct

Ex. We know atoms are divisible into even smaller particles

We know an element can have atoms with different masses

Although john dalton thought atoms were indivisible, investigators in the late 1800s proved otherwise. As scientific advances allowed a deeper exploration of

matter, it became clear that atoms are actually composed of several basic types of smaller particles

and that the number and arrangement of these particles within an atom determine that atom’s

chemical properties. Today we define an atom as the smallest particle of an element that retains the

chemical properties of that element.

The Structure of the Atom

All atoms consist of two regions

1. Nucleus very small region located near the center of an atom

1. In the nucleus there is at least one positively charged particle called the proton

2. Usually at least one neutral particle called the neutron

2. Surrounding the nucleus is a region occupied by negatively charged particles called electrons

Discovery of the Electron

Resulted from investigations into the relationship between electricity and matter

Late 1800s, many experiments were performed electric current was passed through different gases at low pressures

These experiments were carried out in glass tubes known as cathode-ray tubes

Cathode Rays and Electrons Investigators noticed

that when current was passed through a cathode-ray tube, the opposite end of the tube glowed

Hypothesized that the glow was caused by a stream of particles, which they called a cathode ray

The ray traveled from the cathode to the anode when current was passed through the tube

Observations

1. cathode rays deflected by magnetic field in same way as wire carrying electric current (known to have negative charge)

2. rays deflected away from negatively charged object

Observations led to the hypothesis that the particles that compose cathode rays are negatively charged

Strongly supported by a series of experiments carried out in 1897 by the English physicist Joseph John Thomson

He was able to measure the ratio of the charge of cathode-ray particles to their mass

He found that this ratio was always the same, regardless of the metal used to make the cathode or the nature of the gas inside the cathode-ray tube

Thomson concluded that all cathode rays are composed of identical negatively charged particles, which were later named electrons

Charge and Mass of the Electron

Confirmed that the electron carries a negative electric charge

Because cathode rays have identical properties regardless of the element used to produce them, it was concluded that electrons are present in atoms of all elements

Cathode-ray experiments provided evidence that atoms are divisible and that one of the atom’s basic constituents is the negatively charged electron

Thomson’s experiment revealed that the electron has a very large charge for its tiny mass

Mass of the electron is about one two-thousandth the mass of the simplest type of hydrogen atom (the smallest atom known)

Since then found that the electron has a mass of 9.109 × 10−31 kg, or 1/1837 the mass of the hydrogen atom

Based on information about electrons, two inferences made about atomic structure

1. b/c atoms are neutral they must have positive charge to balance negative electrons

2. b/c electrons have very little mass, atoms must have some other particles that make up most of the mass

Thomson’s Atom Plum pudding model (based on English dessert)

Negative electrons spread evenly through positive charge of the rest of the atom

Like seeds in a watermelon

Discovery of Atomic Nucleus

1911 by New Zealander Ernest Rutherford and his associates Hans Geiger and Ernest Marsden

Bombarded a thin, gold foil with fast-moving alpha particles (positively charged particles with about four times the mass of a hydrogen atom)

Assume mass and charge were uniformly distributed throughout atoms of gold foil (from Thomson’s model of the atom)

Expected alpha particles to pass through with only slight deflection

What Really Happened… Most particles passed with only slight

deflection

However, 1/8,000 were found to have a wide deflection

Rutherford explained later it was “as if you have fired a 15-inch artillery shell at a piece of tissue paper and it came back and hit you.”

Explanation After 2 years, Rutherford finally came up with an

explanation

The rebounded alpha particles must have experienced some powerful force within the atom

The source of this force must occupy a very small amount of space because so few of the total number of alpha particles had been affected by it

The force must be caused by a very densely packed bundle of matter with a positive electric charge

Rutherford called this positive bundle of matter the nucleus

Rutherford had discovered that the volume of a nucleus was very small compared with the total volume of an atom

If the nucleus were the size of a marble, then the size of the atom would be about the size of a football field

But where were the electrons?

Rutherford suggested that the electrons surrounded the positively charged nucleus like planets around the sun

He could not explain, however, what kept the electrons in motion around the nucleus

Rutherford’s Atom

Composition of Atomic Nucleus

Except hydrogen, all atomic nuclei made of two kinds of particles Protons Neutrons

Protons = positive

Neutrons = neutral

Electrons = negative

Atoms are electrically neutral, so number of protons and electrons IS ALWAYS THE SAME

The nuclei of atoms of different elements differ in the number of protons they contain and therefore in the amount of positive charge they possess

So the number of protons in an atom’s nucleus determines that atom’s identity

Forces in the Nucleus Usually, particles that have the same electric

charge repel one another

Would expect a nucleus with more than one proton to be unstable

When two protons are extremely close to each other, there is a strong attraction between them

Nuclear forces short-range proton-neutron, proton-proton, and neutron-neutron forces that hold the nuclear particles together

The Sizes of atoms Area occupied by electrons is electron cloud –

cloud of negative charge

Radius of atom is distance from center of nucleus to outer portion of cloud

Unit – picometer (10-12 m)

Atomic radii range from 40-270 pm

Very high densities – 2 x 108 tons/cm3

Counting AtomsConsider Neon, Ne, the gas used in many illuminated signs. Neon is a minor part of the atmosphere. In fact, dry air contains only about 0.002% Ne. And yet there are about 5 x 1017 atoms of neon present in each breath you inhale. In most experiments, atoms are too small to be measured individually. Chemists can analyze atoms quantitatively, however, by knowing fundamental properties of the atoms of each element. In this section, you will be introduced to some of the basic properties of atoms. You will then discover how to use this information to count the number of atoms of an element in a sample with a known mass. You will also become familiar the the mole, a special unit used by chemists to express amounts of particles, such as atoms and molecules.

Atomic Number Atoms of different elements

have different numbers of protons

Atomic number (Z) number of protons in the nucleus of each atom of that element

Shown on periodic table

Atomic number identifies an element

3Li

Lithium6.941

[He]2s1

Isotopes Hydrogen and other atoms

can contain different numbers of neutrons

Isotope atoms of same element that have different masses

Isotopes of Hydrogen

Mass Number

Mass number total number of protons and neutrons in the nucleus of an isotope

Naming Isotopes Isotopes are usually identified by specifying their mass

number

There are two methods for specifying isotopes

1. Mass number is written with a hyphen after the name of the element

1. Tritium is written as hydrogen-3

2. Refer to this method as hyphen notation

2. Shows the composition of a nucleus as the isotope’s nuclear symbol

1. Uranium-235 is written as 23592U

2. The superscript indicates the mass number and the subscript indicates the atomic number

Nuclide – general term for specific isotope of an element

Sample Problem How many protons, electrons, and neutrons are there in an

atom of chlorine-37?

atomic number = number of protons = number of electrons

mass number = number of neutrons + number of protons

mass number of chlorine-37 − atomic number of chlorine = number of neutrons in chlorine-37

mass number − atomic number = 37 (protons plus neutrons) − 17 protons = 20 neutrons

Practice Problems1. How many protons, electrons, and neutrons are in an atom of

bromine-80?

Answer 35 protons, 35 electrons, 45 neutrons

2. Write the nuclear symbol for carbon-13.

Answer 136C

3. Write the hyphen notation for the element that contains 15 electrons and 15 neutrons.

Answer phosphorus-30

Relative Atomic Masses Because atoms are so small, scientists use a standard to

control the units of atomic mass carbon-12

Randomly assigned a mass of exactly 12 atomic mass units (amu)

1 amu is exactly 1/12 the mass of a carbon-12 atom

Atomic mass of any atom determined by comparing it with mass of C-12

Ex. H 1/12 C-12, so 1 amu

Average Atomic Masses

Average atomic mass the weighted average of the atomic masses of the naturally occurring isotopes of an element

Ex. You have a box containing 2 size of marbles 25% are 2.00g each 75% are 3.00g each

25 x 2.00g = 50g 75 x 3.00g = 225g

50g + 225g = 275g

Divide by 100 – average marble mass = 2.75g

Method 1

25 x 2.00g = 50g 75 x 3.00g = 225g

50g + 225g = 275g

Divide by 100 – average marble mass = 2.75g

Method 2

25% = 0.25 75% = 0.75

(2.00g x 0.25) + (3.00g x 0.75) = 2.75g

Calculating Average Atomic Mass

Copper

69.17% Cu-63 – 62.929599 amu

30.83% Cu-65 – 64.927793 amu

(0.6917 x 62.929599) + (0.3083 x 64.927793)

= 63.55 amu

Relating Mass to Numbers of Atoms

The relative atomic mass scale makes it possible to know how many atoms of an element are present in a sample of the element with a measurable mass

Three very important concepts provide the basis for relating masses in grams to numbers of atoms

1. The mole

2. Avogadro’s number

3. Molar mass

The Mole SI unit for an amount of substance (like 1 dozen = 12)

Mole (mol) amount of a substance that contains as many particles as there are atoms in exactly 12g of C-12

Avogadro’s Number Avogadro’s number the number of particles in exactly one mole

of a pure substance

6.022 x 1023

How big is that?

If 5 billion people worked to count the atoms in one mole of an element, and if each person counted continuously at a rate of one atom per second, it would take about 4 million years for all the atoms to be counted

Molar Mass Molar mass the mass of one mole of a pure substance

Written in unit g/mol

Found on periodic table (atomic mass)

Ex. Molar mass of H = 1.008 g/mol

Example

How many grams of helium are there in 2 moles of helium?

2.00 mol He x = ? g He

2.00 mol He x 4.00 g He = 8.00 g He

1 mole He

Practice ProblemsWhat is the mass in grams of 3.50 mol of the element copper, Cu?

222 g Cu

What is the mass in grams of 2.25 mol of the element iron, Fe?

126 g Fe

What is the mass in grams of 0.375 mol of the element potassium, K?

14.7 g K

What is the mass in grams of 0.0135 mol of the element sodium, Na?

0.310 g Na

What is the mass in grams of 16.3 mol of the element nickel, Ni?

957 g Ni

1. A chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced?

0.441 mol Al

2. How many moles of calcium, Ca, are in 5.00 g of calcium?

0.125 mol Ca

3. How many moles of gold,Au, are in 3.60 × 10−10 g of gold?

1.83 × 10−12 mol Au

Conversions with Avogadro’s Number

How many moles of silver, Ag, are in 3.01 x 1023 atoms of silver?

Given: 3.01 × 1023 atoms of Ag

Unknown: amount of Ag in moles

Ag atoms × = moles Ag

3.01x1023atomsAg x =0.500 mol Ag

Practice problems1. How many moles of lead, Pb, are in 1.50 × 1012 atoms of

lead?

2.49 × 10−12 mol Pb

2. How many moles of tin, Sn, are in 2500 atoms of tin?

4.2 × 10−21 mol Sn

3. How many atoms of aluminum, Al, are in 2.75 mol of aluminum?

1.66 × 1024 atoms Al

1. What is the mass in grams of 7.5 × 1015 atoms of nickel, Ni?

7.3 × 10−7 g Ni

2. How many atoms of sulfur, S, are in 4.00 g of sulfur?

7.51 × 1022 atoms S

3. What mass of gold,Au, contains the same number of atoms as 9.0 g of aluminum,Al?

66 g Au