Unit 2: Essentials of Chemistry Chapter 1-2, 4-5

Objectives 8 explain the nature of science including the use of the validity of the scientific method and the

difference between a hypothesis, theory and law

9 explain the three major states of matter and their physical and chemical characteristics

10 explain and give examples of physical properties and chemical properties

11 explain and give examples of physical changes and chemical changes

12 identify the five pieces of evidence that a chemical reaction has occurred

13 define and classify matter including pure substances mixtures.

14 understand the historical progression of structure of the atom including such models created by Democritus, Dalton, Thomson, Rutherford, Bohr, and Schrodinger

15 define and identify examples of the laws of conservation of mass, definite proportions, and multiple proportions

16 identify the structure of an atom including the relative masses of a proton, neutron, and electron, their relative charges, and locations in the atom

17 define and identify isotopes and ions

18 define and distinguish between atomic number and mass number and identify the parts of the nuclear symbol

19 define and calculate average atomic mass in amu’s

20 understand and record the arrangement of electrons in an atom including Hund’s rule, the Pauli exclusion principle, and the Aufbau principle by writing the orbital notation and electon configuration for specific elements

8 Explain the Nature of Science

The nature of science refers to how science is actually performed.

It explains the scientific view point on solving problems and why certain ideas are not considered.

There are several nature of science ideas which will be discussed on the next few slides

Nature of Science Ideas

1. While science strives to be objective, it cannot be entirely.

◦ It is human nature to use previous experiences and knowledge. This allows scientists to make logical conclusions even though it is not completely objective.

◦ For example: Imagine you were trying to determine an object in a shoebox. An objective scientist would have to try everything. This would be tedious and waste time because certain objects would not fit (like an elephant). By using prior knowledge, certain objects can be eliminated and the object can be identified faster.

Nature of Science Ideas

2. Science needs experimental evidence.

◦ In order for something to be considered

scientifically supported, it must be able to be

tested.

◦ Only testable evidence will be accepted by the

scientific community. They will not accept

unless some experiment supports it. Great

scientists have proposed ideas well ahead of

their time, but these ideas were not accepted

because there was no evidence to support

their claims.

Nature of Science Ideas

3. Science neither accepts or denies the existence of the supernatural.

◦ Supernatural beings are considered to be outside the realm of science.

◦ A supernatural being can change the natural world at their will so science cannot test for their existence.

◦ Because science cannot test it, science cannot make a statement one way or the other on the existence of the supernatural.

◦ Science only tries to develop conclusions based on ideas that can be tested.

Nature of Science Ideas

4. Science cannot prove anything 100% which makes it a flexible subject. ◦ As that science relies on data and evidence, it can

only support ideas.

◦ As new evidence is generated, this support can strengthen or weaken.

◦ Often new technology allows for new information to be discovered.

◦ For example: In ancient Greece, it was believed atoms were the smallest particles. However, in the early 1900s, scientists used new technology to show that atoms were made of sub-atomic particles (protons and electrons). The idea that atoms were the smallest particles was no longer supported with this new data.

The Scientific Method

When looking at how science is performed,

the scientific method must be discussed.

It is designed with a few steps which are:

◦ Observe a situation

◦ Make a hypothesis

◦ Test the hypothesis

◦ Make a conclusion

Other sources may provide more steps to

the method but it can be condensed to

these four.

The Scientific Method

The scientific method is an excellent

problem solving guide because it is flexible.

It allows for the scientists to start at any

point and proceed with solving the problem

◦ For example, imagine you are performing an

experiment and observe a change unrelated to

your experiment. This could be the beginning of

a new hypothesis yet it did not come from simply

observing the natural world.

This method for solving problems works for

all problems and not just scientific ones.

Hypothesis, Theory, and Law

These three terms are often confused when discussing the nature of science.

A hypothesis is an educated analysis of an observation. It requires testing and is essential an explanation of a phenomena.

A theory is a hypothesis with enough evidence to be accepted as fact. It explains the how and why qualitatively.

A law is a hypothesis with enough evidence to be accepted as fact. It describes the what quantitatively.

Hypothesis, Theory, and Law

Hypothesis

Theory Law

It is possible to change a theory or a law if

there is evidence to support such a change.

Can become with enough

experimental evidence

Can become with enough

experimental evidence

Both are accepted as fact and

cannot become the other

9 States of Matter

There are actually five known states of

matter. Three are common and will be

discussed in greater detail.

◦ Solid

◦ Liquid

◦ Gas

◦ Plasma (occurs in the sun)

◦ Bose-Einstein Condensate (occurs when

temperatures approach -273°C)

Solids

Solids have a defined shape and a defined

volume.

The atoms that make up its structure are

very close together and are limited in their

movement.

Liquids

Liquids have a undefined shape but a defined

volume.

The atoms that make up its structure are

free to move around within its volume.

Gases

Gases have an undefined shape and an

undefined volume.

Their atoms are free to move about the

entire volume presented.

10. Properties

Each substance has a set of properties

(descriptions). These can fall into two

categories:

Physical and Chemical

Physical properties describe

or measure the object

without changing it.

Chemical properties describe

the substance’s ability to

undergo a change.

Example: color, smell, mass Example: ability to rust, flamability

11. Changes

Each substance also can undergo a change

and the same two categories apply:

Physical and Chemical

Physical change is a change

in which the substance is

altered but can be returned

to its original state.

Chemical change is a change in

which the substance is altered

at the molecular level and

cannot be returned to its

original state.

Example: smashing, tearing Example: baking, burning

Phase changes

Phase changes occur when one state of matter becomes another.

◦ Melting S L

◦ Freezing L S

◦ Vaporization L G

◦ Condensation G L

◦ Sublimation S G

◦ Deposition G S

Since these changes can be reversed to get the original back, phase changes are a type of physical change.

12. Evidence for a chemical change

Chemical changes can be more challenging

to determine.

To help in this matter, there a some pieces

of evidence to look for in the event of a

chemical change.

◦ Color change

◦ Formation of a Precipitant

◦ Emission of heat

◦ Emission of light

◦ Emission of a gas

The videos these are linked

to are from youtube.com as

of July 26, 2011

13. Classifying Matter

As discussed earlier, there are three main

types of matter: solid, liquid, and gas.

Matter can be broken down into two

categories: Pure substances and Mixtures

◦ A pure substance consists of only one

component and it has unique chemical and

physical properties.

◦ A mixture is a combination of two or more pure

substances.

Pure Substances

Pure substances can be broken down into a

few different terms. All of the following have

unique chemical and physical properties.

Term Description

Atom The smallest particle with unique

characteristics.

Element Multiple atoms of the same type

Compound Two or more different atoms bonded

together

Molecule Two or more atoms (can be the same

atom) bonded together

Allotropes

An allotrope is an unique type of molecule.

Allotropes are atoms of the same type that

bond to themselves in multiple ways.

◦ For example: Oxygen

Oxygen gas is O2 which means there are two oxygen

atoms

Ozone gas is O3 which means there are three oxygen

atoms.

Mixtures

Mixtures are two or more pure substances

and can be mixed differently.

◦ Homogenous mixtures are thoroughly mixed and

the parts are uniform throughout.

Any solution is a good example.

◦ Heterogeneous mixtures are not uniform and

the parts can be seen mixed throughout.

Most mixtures are heterogeneous and a good example

of this would be a chocolate chip cookie.

14. History of the atom • ~400 BC Introduced the indestructible atom Democritus

• Late 1700s Law of Conservation of Mass Antoine Lavoisier

• Early 1800s Atomic Theory John Dalton

• 1897 Discovered the electron JJ Thomson

• 1898 Discovered radioactivity Marie Curie

• 1909 Gold-Foil Experiment Ernest Rutherford

• 1913 Determined the charge of an electron Robert Millikan

• 1913 Discovered energy levels for electrons Niels Bohr

• 1926 Discovered atomic orbitals Erwin Schrodinger

• 1932 Discovered the neutron James Chadwick

• 1935 Discovered artificial radioactivity Irene Jolliet-Curie

Dalton’s Atomic Theory

1. Elements are made of tiny particles called atoms.

2. Atoms of an element are different from other elements and can be distinguished by their atomic masses.

3. All atoms of a given element must be identical in properties.

4. Atoms of an element can combine with atoms of a different element to for compounds.

5. Chemical reactions rearrange the atoms but cannot create or destroy atoms. Return

Thomson’s Plum Pudding Model

As that Thomson discovered the

electron, that meant the atom

contained smaller parts.

This would change the model used

for the atom.

◦ The model used to be a solid sphere

◦ The plum pudding model used

electrons as the “plums” and the rest

of the atom as the “pudding”.

◦ The plums were negative and the

pudding was positive.

Return

Gold Foil Experiment

The gold foil experiment was conducted by Ernest Rutherford and his graduate students, Hans Geiger and Robert Marsden.

By shooting alpha particles at a gold foil, they noticed the particles essentially went straight through.

This led them to conclude the atom was mostly empty space with a dense positive core (protons.

Return

Picture was taken from: http://www.kentchemistry.com/links/AtomicStructure/rutherfordtutorial.htm on

July 28, 2011.

Bohr Model of the Atom

Niels Bohr created an atomic model after

doing work with the color spectra

emitted from a hydrogen atom.

His model is sometimes called the solar

system model.

The following link provides a more

detailed description (the first 15 slides

covers Bohr):

http://science.sbcc.edu/physics/solar/science

segment/bohratom.swf Return

Schrodinger’s Orbitals

The current model of the atom uses the orbitals discovered by Schrodinger.

Within each energy level, there exists four kinds of orbitals: s, p, d, and f.

Each can hold a certain number of electrons.

The shape of each orbital is shown on the next slide.

◦ The image was taken from: http://chemwiki.ucdavis.edu/Physical_Chemistry/Quantum_Mechanics/Atomic_Theory/Electrons_in_Atoms/Electronic_Orbitals on July 28th, 2011.

Return

15. Define and Identify Three Laws

Law of Conservation of Mass

◦ The law of conservation of mass states that mass

can be mass can neither be created or destroyed.

◦ This means that the products of a chemical

reaction will have the same mass as the

reactants.

15. Define and Identify Three Laws

Law of Definite Proportions

◦ This law states that any sample of a compound

has the same composition.

◦ This means that water will always be H2O

whether it is found in Iowa or somewhere else.

15. Define and Identify Three Laws

Law of Multiple Proportions

◦ This law states that the mass ratio for one of the

elements in a compound that combines with a

fixed mass of another element can be expressed

in small whole numbers.

Compound % Oxygen % Nitrogen %O ÷ %N Mass Ratio

NO2 69.56 30.44 2.285 4

NO 53.32 46.68 1.142 2

N 2O 36.35 63.65 0.571 1

16. Identify the structure of the

atom The atom is constructed of three basic

subatomic particles.

◦ In the last 20 years, it has been discovered that

quarks make up both protons and neutrons.

Subatomic

Particle

Location Charge Approximate

Mass

Proton Nucleus +1 1 amu

Electron Electron Cloud -1 0.00055 amu

Neutron Nucleus 0 1 amu

Structure of the Atom-Bohr Model

This picture came from http://www.universetoday.com/82128/parts-of-an-atom/ on

August 31, 2011.

17. Isotopes and Ions

Each atom can have different variations.

All atoms are identified by the number of

protons they contain.

◦ Oxygen will always have 8 protons. If you added

an additional proton, the atom would no longer

be oxygen (it would be fluorine).

The number of neutrons and electrons can

vary slightly.

Isotopes-changes in neutrons

Protons contain a positive charge. When an atom becomes large, it contains several protons. That much positive charge in one location is unstable.

Neutrons, which have no charge, act as spacers in between the protons.

An isotope is the name of an atom with different amounts of neutrons.

◦ Oxygen for example has three common isotopes.

Oxygen-16 has 8 neutrons

Oxygen-17 has 9 neutrons

Oxygen-18 has 10 neutrons

Ions-changes in electrons

Electrons fill the orbitals surrounding the

nucleus.

This is known as the electron cloud.

In a neutral atom, each proton has an

electron in the electron cloud.

When an atom becomes charged, the

number of electrons no longer matches the

protons.

◦ Positive charges indicate a loss of an electron.

◦ Negative charges indicate a gain of an electron.

18. Nuclear Symbols

Nuclear symbols are useful for determining

the amount of electrons, protons, and

neutrons in an element.

Mass number: Shows

the number of neutrons

and the number of

protons.

Atomic Number: Shows

the number of protons.

Charge: Indicates the

difference between the

electrons and protons.

Negative charges indicate

more electrons while

positive charges indicate

more protons.

19. Average atomic mass

On the Periodic Table, the mass listed is the

average atomic mass.

This is an average of all the naturally

occurring isotopes of an atom.

Calculating an average for a large amount of

particles can be challenging.

Average Atomic Mass

Carbon has three common isotopes.

◦ Carbon-12

◦ Carbon-13

◦ Carbon-14

These three common isotopes do not come

in equal amounts though:

◦ 98.89% is carbon-12

◦ 1.10% is carbon-13

◦ 0.01% is carbon-14

Average atomic mass

To determine the average atomic mass, the

following formula should be used:

(Atomic Mass x Percent Abundance) + (Atomic Mass x Percent Abundance)

+…..=average atomic mass

So in the case of carbon: (12.00 amu x .9889) + (13.00 amu x 0.0111) + (14.00 amu x 0.0001) =

12.01 amu

20. Arrangement of Electrons

According to Bohr and Schrodinger,

electrons surround the atom in energy levels

that take the form of orbitals.

The Periodic Table is broken into four

sections that correlate to the orbitals

The placement of electrons within these

orbitals follow set rules.

Arrangement of Electrons

Hund’s Rule

◦ Electrons will fill the available orbitals at a certain

energy level before pairing.

Pauli Exclusion Principle

◦ Only two electrons can occupy each orbital and

their spins will be opposite.

Aufbau Principle

◦ The lowest-energy orbitals will be filled first.

Below is a depiction of how the Periodic

Table shows the orbitals (SPDF).

S

F

D

P

Energies of Each Level

The orbitals correspond to different levels

of energies.

In general, each row on the Periodic Table

represents a different level of energy.

It gets more complicated farther down the

Periodic Table.

Energies of Each Level 7p ___ ___ ___

6d ___ ___ ___ ___ ___

5f ___ ___ ___ ___ ___ ___ ___

7s ___

6p ___ ___ ___

5d ___ ___ ___ ___ ___

4f ___ ___ ___ ___ ___ ___ ___

6s ___

5p ___ ___ ___

4d ___ ___ ___ ___ ___

5s ___

4p ___

3d ___ ___ ___ ___ ___

4s ___

3p ___ ___ ___

3s ___

2p ___ ___ ___

2s ___

1s ___

As you progress from the 1s to

the 7p, you increase the amount

of energy. Notice how the d-

sublevel is always after the s-

sublevel of the previous energy

level (example 3d follows 4s).

Notice the f-sublevel follows

the s-sublevel of two energy

levels before it (example: 4f

follows 6s).

Energ

y

2.14 Orbital Notation

To show orbital notation, the three rules

must be followed.

◦ Therefore, start at the lowest energy level.

◦ Designate an electron by drawing an arrow

The arrow indicates the spin

◦ Place one electron in each orbital until they each

have one on that level. Then go back and pair

them.

◦ Only two electrons fit in each orbital.

◦ The arrows must point in opposite directions to

show opposite spins.

Orbital Notation

Look at the element sulfur-32.

First determine the number of electrons

◦ Since sulfur has an atomic number of 16, it has

16 protons, 16 electrons, and 16 neutrons.

3p ___ ___ ___

3s ___

2p ___ ___ ___

2s ___

1s ___

Place electrons in the first energy level and

continue up.

When 2p is hit, fill each orbital and then go

back and pair.

Finish by repeating the process used for 2p

but stop when you hit 16 arrows.

This is the orbital notation.

Electron Configuration

Orbital notation can take up a lot of space.

It does a nice job of giving a visual of the

location of each electron.

Electron configuration is a shorthand

notation for determining the location of the

electrons.

It follows the same rules but is slightly easier

to write.

Electron Configuration

As you read across the Periodic Table, you can pick out the

electron configuration.

The electron configuration for oxygen-16 is:

O:1s2 2s2 2p4

Red represents the s-sublevel and yellow is the p-sublevel.

Each row is an energy level.

Since oxygen has eight electrons, we count eight boxes.

The superscripts on the sublevels indicate the number of

electrons.

Electron Configuration

The same rules apply to the d and f

sublevels.

Example: Gold-196 (79 electrons)

Au: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d9

This concludes the tutorial on

measurements.

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Definitions-Select the word to return to the tutorial

Qualitatively

◦ Refers to a description or observations

Quantitatively

◦ Refers to a measured amount; uses numbers

Objective

◦ Not influenced by personal feelings or

interpretations.