UNIT 1: INTRO AND ATOMIC THEORY

INTRODUCTION TO CHEMISTRY

Chemistry: study of matter and its composition, structure, properties and changes

Main Branches: analytical, biochemical, inorganic, organic, physical, theoretical

SCIENTIFIC RESEARCH

• Scientific Method: a logical series of steps used to solve problems

• Steps• State problem

• Gather information (research)

• Hypothesis

• Conduct controlled experiment(s)

• Analyze data

• Draw conclusions

• Revise and repeat steps as needed

COLLECTING DATA

Why do scientists worldwide use the SIsystem of measurement?

•To improve communication

•Convenience

•Diminish likelihood of errors

SI BASE UNITS USED IN CHEMQuality Unit Name Abbrev Tool to

Measure

Length Meter m Ruler/Meter stick

Mass Kilogram kg Balance

Time Second s, sec Clock/Stopwatch

Temperature Kelvin K Thermometer

Amount of substance Mole mol Not applicable

MORE ABOUT SI

Derived unit: combination of basic SI units

ex: Density (mass/volume), volume (l*w*h), velocity (distance/time), force (kg*m/s)

What is the purpose of SI prefixes?

• easy to convert, represent magnitudes of numbers, based on powers of 10

You must

memorize the

prefixes found in

your notes!!

Please begin

memorizing

them starting tonight

1 Mega (M) 106 base

1 kilo (k) 103 base

1 hecto (h) 102 base

1 deka 101base

Base Meter, liter,

gram, second

101 deci (d) 1 base

102 centi (c) 1 base

103 milli (m) 1 base

106 micro (µ) 1 base

109 nano (n) 1 base

1012 pico (p) 1 base

MEASUREMENTS

What is a measurement?

something that has magnitude (#) and a unit (label), and is compared to a standard

Why are measurements always uncertain?

smallest (last) digits always estimated

There are also errors in instruments and human errors!

ACCURACY VS PRECISION

• Accuracy – how close a measurement is to the accepted value (correct)

• Precision – how close a series of measurements are to each other (consistent)

PERCENT ERROR

• Indicates the accuracy of a measurement

• Memorize this formula!!

% 𝑒𝑟𝑟𝑜𝑟 =𝑒𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 − 𝑎𝑐𝑐𝑒𝑝𝑡𝑒𝑑

𝑎𝑐𝑐𝑒𝑝𝑡𝑒𝑑× 100

What is Jake’s percent error, if he measured the density of a substance to be 6.8 g/mL during his experiment and the accepted value is 7.2 g/mL?

ESTIMATED DIGITS

•When you read a piece of equipment and round the last digit.

•You need to use estimated digits any time you complete a measurement in a lab.

ESTIMATED DIGITS AND A GRADUATED CYLINDER

In mL

1) To what place can the

graduated cylinder be

estimated?

2) What is the measurement with

an estimated digit?

ESTIMATED DIGITS AND A GRADUATED CYLINDER

In mL

1) To what place can the

graduated cylinder be

estimated? Hundredths

2) What is the measurement with

an estimated digit?

~ 20.39 mL

• All pieces of equipment need to include an estimated digit.

• Examples:

*Triple Beam Balance

*Thermometer

*Ruler

*Graduated Cylinder

SIGNIFICANT FIGURES

• All measurements are actually estimates!

• Significant figures indicate the precision of a measurement

• Sigfigs in a measurement include the known digits plus a final estimated digit.

COUNTING SIGNIFICANT FIGURES

Count all numbers EXCEPT:•Leading zeros – 0.0025•Trailing zeros without a decimal point – 2,500

EXAMPLES

1. 214 m2. 500. L3. 820.0 L4. 1.0200x105 kg5. 807,000 kg6. 0.080 s

214 L 3 sigfigs500. L 3 sigfigs820.0 L 4 sigfigs1.0200x105 kg 5sigfigs807,000 kg 3 sigfigs0.080 s 2 sigfigs

CALCULATING WITH SIG FIGS

Multiply/Divide:The number with the fewest sig figs determines the number of sig figs in the answer.

Example:

(13.91g/cm3)(23.3cm3) =

CALCULATING WITH SIG FIGS

Add/subtract:

The number with the lowest decimal value determines the place of the last sig fig in the answer.

Examples:

2.65m+5.3m= 189L+270L=

CALCULATING WITH SIG FIGS

More practice: 1) 19.82 g ÷ 9.1 g

2) 18.9 g – 0.84 g

SCIENTIFIC NOTATION

Why is scientific notation used?

To express very large or very small numbers easily

MORE ON SCI NOTATION

A positive exponent represents a number _greater_ than one.

Ex: 6.8 x 103 = 6,800

A negative exponent represent a number _less_ than one.

Ex: 6.8 x 10-3 = 0.0068

SCIENTIFIC NOTATION EXAMPLES

Write the following in scientific notation

1)5, 600, 000 m

Write the following in standard notation

1)9.12 x 10-3 cg

TYPES OF GRAPHS

• Bar Graph

• Line Graph

• Scatter Plot

DATA PRESENTATION

What are the most common types of graph used in chemistry?

Line graph and Scatter plot

When constructing a graph you should?

• Title graph

• Label axes (x-axis = independent, you control) (y-axis = dependent variable)

• Scale axes evenly (spread out as much as possible)

INTERPRETING RELATIONSHIPS FROM GRAPHS

Direct proportions: both variables change in the same direction

y = x

Example: Temperature vs. Pressure

INTERPRETING RELATIONSHIPS FROM GRAPHS

Inverse proportions: variables change in opposing directions

𝑦 =1

𝑥

Example: Temperature vs. Volume

DIMENSIONAL ANALYSIS

• Dimensional Analysis is a problem solving method that allows us to use conversion factors to change between different units of measurements

EXAMPLE!

Convert 13.4 g to mg

EXAMPLE!

Convert 267 kL to L

EXAMPLE!

Convert 83.12 cm to hm

EXAMPLE!

Convert 72.8 mph to m/sNote: 1 mile = 5280 ft; 1 meter = 3.28 ft

DENSITY

•Density: a ratio that compares the mass of an object to its volume

•𝐷𝑒𝑛𝑠𝑖𝑡𝑦 =𝑀𝑎𝑠𝑠

𝑉𝑜𝑙𝑢𝑚𝑒or 𝐷 =

𝑀

𝑉

Example:

An object has a volume of 825 cm3 and a density of 13.6 g/cm3. Find its mass.

MATTER

•Matter is anything that has mass and takes up space

•Matter can be classified in two groups:•Mixtures

•Pure substances

PURE SUBSTANCES

•Two types:•Elements Ex: Oxygen (O2)

•Compounds Ex: Water (H2O)

•The smallest individual unit of an element is called an atom

•Atoms of two or more elements chemically combined make up a compound.

WHAT MAKES UP AN ATOM?

Particle SymbolRelative

Charge

Relative

Mass

Location in

the Atom

Proton 𝑝+,11𝐻 +1 1 nucleus

Neutron 01𝑛 0 1 nucleus

Electron 𝑒−,−10𝑒 -1 0

Electron

cloud

outside the

nucleus

QUARKS•The atom consists of three subatomic particles, the proton, neutron, and electron. The electron is an elemental particle, whereas the proton and neutron consist of three smaller particles called quarks.

DEFINITIONS

•Nucleons: particles in the nucleus (protons and neutrons)

•Atomic number: number of protons in the nucleus, identifies the element, “Z”

•Mass number: number of protons and neutrons in an element, “A”, (NOT found on the periodic table)

DEFINITIONS CONT.

• Isotopes: atoms of the same element with different numbers of neutrons

DEFINITIONS CONT.

•Atomic number: average atomic mass of an atom, found on the periodic table

DEFINITIONS CONT.

•Nuclides: a particular isotope of an element

•Examples: C-12, C-14, 36𝐿𝑖

• Ion: an atom that has lost or gained an electron•Lost electron → positive charge

•Gained electron → negative charge

𝟏𝟑𝟐𝟕𝑨𝒍𝟑+

Charge

ATOMIC MASS

•The average mass system is a relative system based on the mass of a standard nuclide.

•This nuclide is carbon-12.

•The unit for measuring atomic masses is the atomic mass unit (amu)

CALCULATE ATOMIC MASS

•Use the following formula:

(% 𝒊𝒏 𝒅𝒆𝒄𝒊𝒎𝒂𝒍 𝒇𝒐𝒓𝒎)(𝒎𝒂𝒔𝒔 𝒐𝒇 𝒏𝒖𝒄𝒍𝒊𝒅𝒆)

CALCULATE ATOMIC MASS

•Calculate the average atomic mass of boron if a sample contains• 19.78% boron-10 (atomic mass=10.013amu)

• 80.22% boron-11 (atomic mass=11.009amu)

THE HISTORY OF THE ATOMHow did we learn about the atom?

THE GREEK PHILOSOPHERS, 400 B.C.• Democritus proposed the first

idea of the atom around 420 BC

• Democritus coined the term, “atom”

• Derived from the Greek word atomos, meaning “indivisible.”

• Are atoms indivisible?

THE 1ST SCIENTIFIC THEORY, 1803

• Dalton’s Atomic Theory

1. All matter is made up of indestructible atoms.

2. Atoms of the same element are identical.

3. Atoms of one element are different from atoms of another element.

4. Atoms can physically mix together or can chemically combine in whole-number ratios to form compounds.

Still True Today?

1. Yes, but not “indestructable”

2. No, isotopes!

3. Yes: differ by # protons

4. Yes

DALTON BASED HIS THEORY OF THE WORK OF TWO SCIENTISTS

• Lavoisier (1788)• Law of Conservation of

Mass:

• Matter cannot be created or destroyed only rearranged.

• Proust (1794)• Law of Definite

Proportions

• A pure substance always contains the same elements combined in the same ratio by mass, Ex: H2O always 2:1

LAW OF MULTIPLE PROPORTIONS

• 1810

• Dalton used his theory to develop this law

• The same two elements can combine in different ratios and form different compounds

• Ex: CuCl and CuCl2; CO and CO2

1ST DISCOVERY OF SUBATOMIC PARTICLES, 1897

•J. J. Thomson•He used cathode rays to discover the electron. He measured the bending of the rays to determine the mass to charge ratio of the electron.

HOW DID THOMSON DISCOVER THE ELECTRON?

• Passed an electric current through gases in a cathode-ray tubeproducing a glowing beam.

• If electrically charged plates are placed near the cathode tube, the cathode ray will be attracted towards the positive plate and repelled from the negative plate.

• Therefore, the ray must be made of negatively-charged particles.

CATHODE RAY TUBES CONT.

•Modified cathode ray tubes were used to discover a beam of positive charge. These rays were determined to be positive particles and were named protons.

1ST DISCOVERY OF SUBATOMIC PARTICLES, 1897

Thomson developed a new model called the plum pudding model.

Plum Pudding Model of the Atom: Tiny particles of negative charge embedded in a ball of positive charge.

MILLIKAN’S OIL DROP EXPERIMENT, 1916

• Found the mass and charge of the electron.

-

Force

upwards

from the

negatively

charged

plate.

Force of Gravity

HERE’S WHAT MILLIKAN’S APPARATUS REALLY LOOKED LIKE…• He sprayed oil mist into

chamber. Electrons transferred

to droplets

• Oil drops fall through the

chamber due to gravity

• Change of plants adjusted to

offset gravity and suspend

droplet.

• Charge of electron calculated.

CHADWICK, 1932

• Discovered the neutron, MUCH LATER!

• Beryllium bombarded with alpha particlesto form carbon and high energy neutral particle (neutron)

THE 1ST SCIENTIFIC THEORY, 1803

• Dalton’s Atomic Theory

1. All matter is made up of indestructible atoms.

2. Atoms of the same element are identical.

3. Atoms of one element are different from atoms of another element.

4. Atoms can physically mix together or can chemically combine in whole-number ratios to form compounds.

Still True Today?

1. Yes, but not “indestructable”

2. No, isotopes!

3. Yes: differ by # protons

4. Yes

OKAY, SO NOW WE HAVE SUBATOMIC PARTICLES…

How are the particles arranged in the atom???

RUTHERFORD’S GOLD FOIL EXPERIMENT, 1911

• The Nuclear Model of the Atom –

The nucleus was the next major

focus in the development of the

atomic model.

• Ernest Rutherford studied the

radiation emitted by these

substances, especially the alpha

particles.

HERE’S WHAT YOU NEED TO REMEMBER:

Rutherford shot alpha particles (2+ charge) at gold foil.

1) Most of the alpha particles went straight through… => Atoms are mostly empty space.

2) A few bounced back… => Small, dense positively-charged nucleus.

Rutherford’s model is called the nuclear model of the atom.

BOHR’S PLANETARY MODEL, 1913

• Electrons orbit the nucleus like planets orbit the sun.