atoms

Physical Science chapter 10 1

Atoms

Physical Science

Chapter 4


The atom The atom – smallest piece of matter that has

the properties of an element. Made of

Protons Neutrons Electrons

If we split an atom, we no longer have a specific element We can’t tell an oxygen proton apart from a

carbon proton


Early atomic theory - Democritus Greek philosopher about 400 B.C. Gave us the word atom

Atomos - indivisible. Thought

The world was made of empty space and particles called atoms.

There were different types of atoms for different types of materials.

Theory was not supported by experimental evidence.


Early atomic theory – Aristotle Aristotle did not believe in atoms

thought matter was continuous He was very influential, so Democritus’s

theory was not accepted for many centuries.


17th century

People began to express doubts in Aristotle’s theory.

Experiments were being used to determine the validity of a theory.


John Dalton – early 1800s

Studied experimental observations of chemical reactions

Proposed explanation of these experimental results


Dalton’s Hypothesis1. All matter is composed of very small particles

called atoms.

2. All atoms of an element are exactly alike; atoms of different elements are very different.

3. Atoms cannot be subdivided, created, or destroyed.

4. Atoms unite with other atoms in simple ratios to form compounds

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


Dalton’s theories

However, there some major exceptions to the rules. Some elements can combine with each other in

different proportions H2O and H2O2

Splitting atoms Different atoms of the same element


Cathode rays and electrons

1897 – J.J. Thomson tested cathode rays and discovered that they were electrons. they had mass they were negatively charged


Thomson’s plum pudding model

In this model, the raisins were the electrons and the pudding was the positive charge.

Sort of like chocolate chip cookie dough. The chips are the electrons and the dough is the

positive charge. Explained the experiments that had been

done so far.


Testing the plum pudding model See page 117 fired alpha particles at a very thin (a few

atoms thick) sheet of gold foil. They expected the particles to go right

through because the spread out positive charge in the “pudding” wouldn’t be strong enough to deflect them.


What happened

Most of the particles did go right through without being deflected at all.

Some were deflected at large angles. Ernest Rutherford explained it:

the positive charge on the atom was concentrated at a small core – now called the nucleus.


The atom as we now “know” it The nucleus contains all of the positive

charge and most of the mass. The negatively charged electrons have very

small mass and are located around the nucleus in the electron cloud.

Most of an atom is empty space.


Discuss

1. Compare and contrast Thomson’s atomic model with Rutherford’s atomic model.

2. How did the gold foil experiment lead to the conclusion that the atom has a nucleus?


Atoms

Basic building blocks of matter Smallest part that can be called an element

Made up of: Nucleus – in the center

Protons – positively charged Neutrons – neutral (no charge)

Electrons – around the nucleus and negatively charged


Mass and charge comparisons A proton’s positive charge is equal to an

electron’s negative charge. A proton and a neutron have about the same

mass. An electron has a mass that is about 1/2000

the mass of a proton Electrons are much smaller! They are so small

that their mass is negligible.


Atomic number (Z)

Given on the periodic table. The atomic number of an atom is the number

of protons it has. Defines what element an atom is

It is also the number of electrons the atom has.

Since an atom has an equal number of positive protons and negative electrons, the whole atom is electrically neutral.


Mass number (A)

the number of particles in the nucleus of an atom.

In other words, the sum of the number of protons and the number of neutrons.


Calculating the number of neutrons

To find the number of neutrons, just subtract the atomic number from the mass number.

number atomic number mass neutrons ofnumber -


Isotopes

Atoms of the same element that have different numbers of neutrons.

All the isotopes of an atom have the same number of protons and electrons, it is only the number of neutrons that is different.

All the isotopes of an atom are chemically the same, even though they differ in mass.


Different isotopes

To distinguish between the two types of neon isotopes, we can write neon-20 and neon-21.


Discuss

Read the “Why It Matters” section on page 123.

Discuss it with your groups – including the critical thinking question.

Physical Science chapter 10 23October 09

Atomic mass

A proton or a neutron has a mass of about 1.7 x 10-27 kg.

This is a very small number, so it is not very convenient to write or work with.

Instead, we use atomic mass units (amu). Also called unified atomic mass units (u)


Atomic mass units

1 amu is defined as one-twelfth the mass of a carbon atom with 6 protons and 6 neutrons.

Since the mass of electrons are negligible, a proton or a neutron has a mass of about 1 amu.


Atomic mass

On the periodic table, the atomic masses listed are usually not whole numbers.

This is because atoms have different isotopes.

So, the masses given on the table are average masses.


Example

Neon has isotopes with mass numbers of 20 and 21.

The atomic mass on the periodic table is given as 20.179.

This tells us that most neon isotopes have a mass number of 20.

How many neutrons are in each neon isotope?


Weighted averages

We then use a weighted average to find the average mass of an atom of a given element.

This is called the average atomic mass or just atomic mass.


The Mole

SI unit for amount of substance Abbreviated mol A counting unit 6.022 x 1023 particles

Avogadro’s number Based on carbon-12, 12 g of C-12 contains

1 mol of atoms


Molar mass

The mass of 1 mol of a pure substance Units: g/mol Numerically equal to the atomic mass in amu

On the periodic table the number with a decimal is the atomic mass in amu AND the molar mass in g/mol


conversions

Grams to moles or moles to grams Use the molar mass


Example

What is the mass in grams of 5.60 mol of sulfur?


Example

How many moles of carbon are in a sample with a mass of 567 g?


You try

How many moles are in 0.255 g of zinc?


You try

What is the mass of 1.21 mol of helium?


Compounds

Also have molar mass Example: What is the molar mass of carbon

dioxide, CO2?

You try: What is the molar mass of hydrogen peroxide, H2O2?


Bohr model

Neils Bohr – 1913 Electrons orbit the nucleus like planets orbit

the sun


Energy Levels

Electrons near the nucleus have low energy. Electrons farther from the nucleus have

higher energy. The electrons are in levels – like floors in a

building. See page 128

Electrons can be on any level, but they cannot be between levels.


Energy levels

The levels are not equally spaced, and they cannot all hold the same number of electrons.

The lowest level can only hold 2 electrons. The second level can hold 8 electrons. The third level can hold 18 electrons.


Valence electrons

Electrons in the outer energy level of an atom Determine the chemical properties of an atom


Electron cloud model

Represents the probable locations of electrons within an atom.

We never know for sure exactly where an electron is because it is so small and it moves so fast.


Orbitals

Spaces in each energy level than electrons occupy

Each orbital can hold two electrons


S orbital

Spherical Each energy level has one Can hold 2 electrons


P orbital

Dumbbell shaped Found in level 2 and up Can have three orientations in space Each p orbital can hold 2 electrons

Total of 6 electrons can be in the p orbitals in an energy level


d and f orbitals More complex shapes There are 5 d orbitals and 7 f orbitals d start in level 3 f start in level 4


Electron Energy States

Ground state – Lowest energy state of an electron

Excited state – higher energy state than the ground state


Electron Transitions Energy gain

Photon absorbed Electron “rides up the elevator” to a higher energy

state Energy loss

Photon emitted Electron “rides down the elevator” to a lower

energy level The amount of energy in the photon

determines how many levels the electron moves


Spectral Analysis

Each element has a unique atomic structure and a unique set of energy levels. Therefore, different atoms emit photons of

different energy. Photons of different energy have different

wavelengths (colors). An element can be identified by the colors it

emits as its electrons lose energy.


Sample Spectra


Discuss

1. State two key features of the modern model of the atom.

2. Describe what happens when an electron jumps from one energy level to another.

atoms

Documents