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PowerPoint® Lecture Slide Presentation
by Patty Bostwick-Taylor,
Florence-Darlington Technical College
Basic
Chemistry
2
PART A
Copyright © 2009 Pearson Education, Inc., publishing as Benjamin Cummings
Chp. 2 Vocabulary
1. Enzyme
2. Cholesterol
3. Glucose
4. Organic Compound
5. Inorganic Compound
6. Decomposition
7. Electron
8. Atomic #
9. Atomic Mass
10.Valence Shell
11.Elements
12.Atoms
13.Atomic Symbol
14.Proton
15.Neutron
16. Molecule
17. Compound
18. Ionic Bond
19. Covalent Bond
20. Hydrogen Bond
21. SynthesisReaction
22. Matter
23. Energy
24. Kinetic Energy
25. Potential Energy
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Matter and Energy
 Matter—anything that occupies space and has
mass (weight)
 Energy—the ability to do work
 Chemical
 Electrical
 Mechanical
 Radiant
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Chemical Energy
 Stored in the bonds of chemical substances
 Broken bonds release (potential) energy and it
becomes kinetic energy
 Ex: gasoline molecules power you vehicle
 Ex: Chemical energy from foods we eat power our
body
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Electrical Energy
 Results from movement of charged particles
 Ex: electricity is from electrons flowing across
wiring
 Ex: Nervous system uses eletrical currents called
nerve impulses to transmit messages throughout
the body.
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Mechanical Energy
 Energy directly involved in moving matter
 Ex: riding a bike; your legs provide the mechanical
energy that moves the pedals
 Ex: muscles in you legs shorten, they pull on your
bones, causing your limbs to move (allowing you
to pedal the bike)
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Radiant Energy
 Travels in waves
 Ex: x-rays, infrared radiation, visible light, radio,
ultraviolet waves, Sonar
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Composition of Matter
 Elements—fundamental units of matter
 96% of the body is made from four elements
 Carbon (C)
 Oxygen (O)
 Hydrogen (H)
 Nitrogen (N)
 Atoms—building blocks of elements
 Hydrogen – simplest atom
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Elements
 Unique substances that cannot be broken down
into smaller substances by ordinary methods
 112 elements are know with certainty; 113-118 are
alleged
 92 of these elements occur naturally
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Atoms
 Building blocks of elements
 Comes from Greek word meaning “incapable of
being divided”
 3 subatomic particles
 Protons – positive charge;
 Neutrons – uncharged
(Protons & electrons are heavy particles and have about the
same mass)
 Electrons – negative charge; (equal to the
positive charge of protons)
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Atomic Structure
 Nucleus
 Protons (p+)
 Neutrons (n0)
 Outside of nucleus
 Electrons (e-)
Figure 2.1
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Identifying Elements
 Atomic number—equal to the number of protons
that the atom contains
 Protons = # of electrons
 Ex: An atom has 5 neutrons, 4 protons, and 4
electrons. Whats the atomic #.
 Atomic mass number—sum of the protons and
neutrons
 Ex: An atom has 5 neutrons, 4 protons, and 4
electrons. Whats the atomic mass #.
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Planetary Model vs. Orbital model
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Atomic Structure of Smallest Atoms
Figure 2.2
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Isotopes and Atomic Weight
 Isotopes
 Have the same number of protons
 Vary in number of neutrons
Figure 2.3
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Isotopes and Atomic Weight
 Atomic weight
 Close to mass number of most abundant
isotope
 Atomic weight reflects natural isotope
variation
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Radioactivity
 Radioisotope
 Heavy isotope
 Tends to be unstable
 Decomposes to more stable isotope
 Radioactivity—process of spontaneous atomic
decay
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Molecules and Compounds
 Molecule—two or more like atoms combined
chemically
 Compound—two or more different atoms
combined chemically
Figure 2.4
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Chemical Reactions
 Atoms are united by chemical bonds
 Atoms dissociate from other atoms when
chemical bonds are broken
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Electrons and Bonding
 Electrons occupy energy levels called electron
shells
 Electrons closest to the nucleus are most
strongly attracted
 Each shell has distinct properties
 The number of electrons has an upper limit
 Shells closest to the nucleus fill first
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Electrons and Bonding
 Bonding involves interactions between electrons
in the outer shell (valence shell)
 Full valence shells do not form bonds
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Inert Elements
 Atoms are stable (inert) when the outermost shell
is complete
 How to fill the atom’s shells
 Shell 1 can hold a maximum of 2 electrons
 Shell 2 can hold a maximum of 8 electrons
 Shell 3 can hold a maximum of 18 electrons
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Inert Elements
 Atoms will gain, lose, or share electrons to
complete their outermost orbitals and reach a
stable state
 Rule of eights
 Atoms are considered stable when their
outermost orbital has 8 electrons
 The exception to this rule of eights is Shell 1,
which can only hold 2 electrons
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Inert Elements
Figure 2.5a
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Reactive Elements
 Valence shells are not full and are unstable
 Tend to gain, lose, or share electrons
 Allow for bond formation, which produces
stable valence
Figure 2.5b
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Chemical Bonds
 Ionic bonds
 Form when electrons are completely
transferred from one atom to another
 Ions
 Charged particles
 Anions are negative
 Cations are positive
 Either donate or accept electrons
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Ionic Bonds
Na
Cl
Sodium atom (Na)
(11p+; 12n0; 11e–)
Chlorine atom (Cl)
(17p+; 18n0; 17e–)
+
–
Na
Cl
Sodium ion (Na+)
Chloride ion (Cl–)
Sodium chloride (NaCl)
Figure 2.6
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Ionic Bonds
Na
Cl
Sodium atom (Na)
(11p+; 12n0; 11e–)
Chlorine atom (Cl)
(17p+; 18n0; 17e–)
Figure 2.6, step 1
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Ionic Bonds
Na
Cl
Sodium atom (Na)
(11p+; 12n0; 11e–)
Chlorine atom (Cl)
(17p+; 18n0; 17e–)
Figure 2.6, step 2
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Ionic Bonds
Na
Cl
Sodium atom (Na)
(11p+; 12n0; 11e–)
Chlorine atom (Cl)
(17p+; 18n0; 17e–)
+
–
Na
Cl
Sodium ion (Na+)
Chloride ion (Cl–)
Sodium chloride (NaCl)
Figure 2.6, step 3
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Chemical Bonds
 Covalent bonds
 Atoms become stable through shared
electrons
 Single covalent bonds share one pair of
electrons
 Double covalent bonds share two pairs of
electrons
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Examples of Covalent Bonds
Figure 2.7a
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Examples of Covalent Bonds
Figure 2.7b
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Examples of Covalent Bonds
Figure 2.7c
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Polarity
 Covalently bonded
molecules
 Some are non-polar
 Electrically
neutral as a
molecule
 Some are polar
 Have a positive
and negative
side
Figure 2.8
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Chemical Bonds
 Hydrogen bonds
 Weak chemical bonds
 Hydrogen is attracted to the negative portion
of polar molecule
 Provides attraction between molecules
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Hydrogen Bonds
Figure 2.9
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Patterns of Chemical Reactions
 Synthesis reaction (A + BAB)
 Atoms or molecules combine to form larger
ones
 Energy is absorbed for bond formation (energy
absorbing reactions)
 Decomposition reaction (ABA + B)
 Molecule is broken down
 Chemical energy is released
 Ex: digestion of food
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Synthesis and Decomposition Reactions
Figure 2.10a
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Synthesis and Decomposition Reactions
Figure 2.10b
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Patterns of Chemical Reactions
 Exchange reaction (AB + CAC + B)
 Involves both synthesis and decomposition
reactions
 Switch is made between molecule parts and
different molecules are made
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Patterns of Chemical Reactions
Figure 2.10c
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