Ch. 1—Chemistry: An Introduction What is chemistry?

• Chemistry is the study of the ___________________ of substances and the changes they undergo.

• It began from “_______________”... the attempts of alchemists to change common metals into _________ through trial and error.

• Do you believe we can make gold? Why or why not?

composition

alchemygold

Half of chemistry in one sentence: “Atoms that don’t have enough electrons in the outer level will fight, barter, beg, make and break

alliances, or do whatever they must to get the right number.”- Kean, Sam. The Disappearing Spoon. New York: Back Bay Publishing, 2010. Print.

How do we classify materials in chemistry?

• Elements cannot be ___________ down or _____________ into simpler substances by chemical means. Elements are the _________ forms of matter that can exists in normal laboratory conditions.

• Compounds are made up of ____ or ________ different elements ______________ bonded together. Compounds can only be broken down into simpler substances by ____________ ____________.

• Mixtures are a physical blend of two or more substances mixed together.” The parts can be separated by _____________ means or ____________ changes.

physicalphysical

broken changed

simplest

2 morechemically

chemical reactions

Chemical Symbols

• Chemists use chemical symbols for the elements involved in a chemical reaction. The symbols are a shorthand way of representing the ______________. (See the Periodic Table for a list of all the symbols.)

• The first letter of the chemical symbol for an element is always _________________.

• The next letter, if needed, is _______________. Each capital letter in a formula, therefore, represents another element.

Examples: ____, ____, Hg, ___, NaBr, ________, LiC2H3O2

• Some symbols come from _______ names: Au=Aurum (Gold)

elements

capitalized

lowercase

H Ne S

Latin

H2O

Chemical Reactions

• When writing chemical reactions, the substances that ___________ with each other are written on the _______ and are called “reactants”.

• The substances that are ____________ are written on the _______ and are called the “products.”

Reactants Products

• The “ ” symbol can be read as “_______” or “reacts to produce.”

Example: 2H2 + O2 2H2O

which means “____________________________________

________________________________________________.”

reactleft

produced right

yields

two hydrogen molecules plus one oxygen

molecule yields two water molecules

Conservation of Mass

• During chemical (or physical) reactions, mass (or matter) is neither _____________ nor _________________.

• The mass of all the reactants _________ the mass of all the products.

• The ___________ of each kind of atom is the same.

• Sometimes it appears that the reactant and product masses are not equal, but a _______ was probably a reactant or product in the reaction, and that is making the difference!

Example: 2H2 + O2 2H2O

• If 4 grams of hydrogen reacted with oxygen to produce 36 grams of water, how many grams of oxygen were used? _______

• Notice that the ____ of H’s and O’s on each side is __________!

created destroyed

equals

number

gas

32

# constant

Conservation of Mass

CaCl2 + Na2SO4 CaSO4 + 2NaCl

mass before = mass after

# atoms before = # atoms after

Atomic Theory and Structure• The smallest particle of an ________________ is an atom.

• The atom is made up of three ________________ particles.The Theory of the Atom

(1) ________________, a famous Greek teacher who lived in the 4th Century B.C., first suggested the idea of the atom.

(2) ________ __________ came up with his solid sphere atomic theory based on the results of his experiments.

(3)The proton has a ______ charge, and it was discovered in _________ by E. Goldstein.

(4)The electron was discovered in _______ by J. J. Thomson by using a cathode ray tube. The electron has a _______ charge. It’s mass is much smaller than the other 2 subatomic particles, therefore it’s mass is usually ______________.

Democritus

John Dalton

element

subatomic

1897(−)

ignored

(+)1886

Cathode Ray Tube

(4) Model:

• a ball of (+) charge containing a number of e-

• no ________________

• often described as the “________ _______________” atom.

Thomson

nucleus

plum pudding

Atomic Models

The Nucleus• (5) Discovered by Ernest ________________ in ________.• He shot a beam of positively charged “alpha particles”,

which are ___________ nuclei, at a thin sheet of ____ ____.

•

• 99.9% of the particles went right on through to the ______________.• Some were slightly deflected. Some even ____________ ________ towards the source! • This would be like shooting a cannon ball at a piece of tissue paper and having it bounce off.

Rutherford 1911

helium gold foil

detector

bounced back

Rutherford’s Experiment

• Most of the atom is more or less _________ ___________. • The nucleus is very _________. (Stadium Analogy)• The nucleus is very _________. (Large Mass ÷ Small Volume)• The nucleus is ______________ charged.

(5) Model: • a ____________ of (+) charge surrounded by

a number of e- • no _____________ and no e- orbitals

Rutherfordnucleus

neutrons

Conclusions about the Nucleus

empty space

densetiny

positively

Nuclear Atomic Structure

• The atom is made up of 2 parts/sections:(1) The ______________ --- (in the center of the atom)(2) The ____________ _________ --- (surrounds the nucleus)

nucleuselectron cloud

(p+ & n0)

e− cloud

(6) The neutron does not have a charge. In other words, it is ________. It was discovered in ____ by James Chadwick. The neutron has about the same _________ as the proton.

• These three particles make up all the ____________________ in the Universe!

• There are other particles such as neutinos, positrons, and quarks, but are typically left for 2nd year chemistry courses.

neutral 1932mass

visible matter

Atomic Models

(7) Model:

• a nucleus of (+) charge that also contains ______________

• nucleus is encircled by e-’s located in definite orbits (or paths).

• e-’s have ___________ energies in these orbits

• e-’s do not lose energy as they orbit the nucleus

Atomic Models

Bohr

neutrons

fixed

Bohr Atomic Model

Bohr Atomic Model

How to draw your own Bohr model

• The atomic number tells you how many electrons a neutral atom will contain.

• The first energy level can only fit _____electrons (like He).

• Each energy level beyond the first one can fit ______ electrons. (or 18 if it is after the middle block which is also called the d-block)

• 2, 8, 8, 18, 18…• We will expand on this when we get to electron configurations.

eight

Quantum Mechanical Model

(8) Mechanical Model ( Wave Mechanical Model)

• no definite ____________ to the e- path (“fuzzy” cloud)

• orbits of e-’s based on the _________________ of finding the e- in the particular orbital shape.

Quantum

shape

probability

Quantum Mechanical Model

Schroedinger's Cat

Quantum Mechanical Model(present)

Counting Subatomic Particles in an Atom• The atomic # of an element equals the number of ____________ in the nucleus. (This number is the whole number on the periodic table)

• The mass # of an element equals the sum of the _____________ and ______________ in the nucleus. (This number cannot be found on the periodic table)

• In a neutral atom, the # of protons = # of ______________. • To calculate the # of neutrons in the nucleus, ______________ the ___________ # from the __________ #.

protons

protonsneutrons

electronssubtract

atomic mass

Practice Problems

(1) Find the # of e-, p+ and n0 for sodium. (mass # = 23)

2) Find the # of e-, p+ and n0 for uranium. (mass # = 238)

3) What is the atomic # and mass # for the following atom? # e- = 15; # n0 = 16

Atomic # = 11 = # e- = # p+ # neutrons = 23-11 = 12

Atomic # = 92 = # e- = # p+ # neutrons = 238-92 = 146

Atomic # = 15 = # e- = # p+ Mass # = p+ + n0 = 15+16 =31

The element is phosphorus!

Isotopes

• An isotope refers to atoms that have the same # of ___________, but they have a different # of ___________.

• Because of this, they have different _________ #’s (or simply, different ___________.)

• Isotopes are the same element, but the atoms weigh a different amount because of the # of ______________.

Examples---> (1) Carbon-12 & Carbon-13 (2) Chlorine-35 & Chlorine-37

(The # shown after the name is the mass #.)• For each example, the elements have identical ___________ #’s,

(# of p+) but different _________ #’s, (# of n0).• Another way to write the isotopes in shorthand is as follows:

C Cl126

3517

The top number is the ________ #, and the bottom # is the __________ number. Calculating the # n0 can be found by _____________ the #’s!

protons neutrons

massmasses

neutrons

atomic mass

mass atomic subtracting

Figure 3.10: Two isotopes of sodium.

More Practice Problems

(1) Find the # e-, p+ and n0 for Xe-131.

2) Find the # e-, p+ and n0 for .

3) Write a shorthand way to represent the following isotope:

# e- = 1 # n0 = 0 # p+ = 1

Cu6329

Atomic # = 54 = p+ = e− n0 = 131 − 54 = 77

Atomic # = 29 = p+ = e− n0 = 63 − 29 = 34

Atomic # = p+ = e− = 1 mass # = n0 + p+ = 1+ 0 = 1

H-1 or H11

Ions

• An atom can gain or lose electrons to become electrically charged.

• Cation = (___) charged atom created by ___________ e-’s.– Cations are ______________ than the original atom.– _____________ generally form cations.

• Anion = (___) charged atom created by _____________ e-’s.– Anions are ____________ than the original atom.– _______________ generally form anions.

Practice Problems: Count the # of protons & electrons in each ion.

a) Mg+2 p+ = _____ e− = ______

b) F−1 p+ = _____ e− = ______

+ losing

smaller

Metals

− gaininglarger

Nonmetals

12 10

9 10

Atomic Mass• Based on the relative mass of Carbon-12 which is exactly

_______.

• 1 p+ ≈ __ atomic mass unit (amu) 1 n0 ≈ __ amu 1e- ≈ __ amu• The atomic masses listed in the Periodic Table are a “weighted

average” of all the isotopes of the element.

12

1 1 0

Weighted AveragePractice Problems: (1) In chemistry semester grades are calculated using a weighted average

of three category scores: Summative Assessments = 80% of your gradeFormative Assessments = 10% of your grade

Semester Exam = 10% of your grade• If a student had the following scores, what would they receive for the

semester? Summative Assessments = 3.5Formative Assessments = 2.5

Semester Exam = 2.0

Weighted AverageStep (1): Multiply each score by the % that it is weighted.Step (2): Add these products up, and that is the weighted average! 3.5 x .80 = 2.80 2.5 x .10 = 0.25

2.0 x .10 = 0.20 Add them up!!

A “normal average” would be calculated by simply adding the raw scores together and dividing by 3…

3.5 + 2.5 + 2.0 = 8 ÷ 3 = 2.6 (C)

+

3.25 (B+)