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Enmanuel’s Chemistry Notebook


Unit 3

Chapter 25 Nuclear Chemistry

The students will learn what happens when an unstable

nucleus decays and how nuclear chemistry affects their lives.

Explore the theory of electromagnetism by comparing and contrasting the

different parts of the electromagnetic spectrum in terms of wavelength,

frequency, and energy, and relate them to phenomena and applications.

Students will be able to compare and contrast the different parts of the

electromagnetic spectrum.

Students will be able to apply knowledge of the EMS to real world phenomena.

Students will be able to quantitatively compare the relationship between energy,

wavelength, and frequency of the EMS.

amplitude

wavelength

frequency

hertz

electromagnetic radiation

photon

Planck’s constant

Explain and compare nuclear reactions (radioactive decay, fission and

fusion), the energy changes associated with them and their associated

safety issues.

Students will be able to compare and contrast fission and fusion reactions.

Students will be able to complete nuclear decay equations to identify the type of

decay.

Students will participate in activities to calculate half-life.

radioactivity

nuclear radiation

alpha particle

beta particle

gamma ray

positron

½ life

transmutation

fission

fusion


Chapter 7

Ionic and Metallic Bonding

The students will learn how ionic compounds form and how

metallic bounding affects the properties of metals.

Compare the magnitude and range of the four fundamental forces

(gravitational, electromagnetic, weak nuclear, strong nuclear).

Students will compare/contrast the characteristics of each fundamental force.

gravity

electromagnetic

strong

weak

Distinguish between bonding forces holding compounds together and other

attractive forces, including hydrogen bonding and van der Waals forces.

Students will be able to compare/contrast traits of ionic and covalent bonds.

Students will be able to compare/contrast basic attractive forces between

molecules.

Students will be able to predict the type of bond or attractive force between

atoms or molecules.

ionic bond

covalent bond

metallic bond

polar covalent bond

hydrogen bond

van der Waals forces

London dispersion forces

Chapter 8

Covalent Bonding

The students will learn how molecular bonding is different

than ionic bonding and electrons affect the shape of a

molecule and its properties.

Interpret formula representations of molecules and compounds in terms of

composition and structure.

Students will be able to interpret chemical formulas in terms of # of atoms.

Students will be able to differentiate between ionic and molecular compounds.

Students will be able to list various VSEPR shapes and identify examples of

each.

Students will be able to predict shapes of various compounds.

Molecule

empirical formula

Atom

Electron

Element

Compound


Enmanuel Garrido

Name ____________________

Go to the web site www.darvill.clara.net/emag

1. Click on “How the waves fit into the spectrum” and fill in this table:

>: look out for the

RED words on the web site!

Low __________, frequency Long wavelength

High frequency, Short ______________

wavelength

Radio Waves

Microwaves Infra-red Visible Light Ultra-violet X-rays

Gamma rays

2. Click on “Radio waves”. They are used for _______________________

communications

3. Click on “Microwaves”. They are used for cooking, mobile _________, Wifi speed _______ cameras cameras and _________. radar

4. Click on “Infra-red”. These waves are given off by _____ hot _________. objects They are used for remote controls,

cameras in police ____________ helicopters , and alarm systems.

5. Click on “Visible Light”. This is used in DVD ___ players and _______ laser printers, and for seeing where we’re going.

6. “UV” stands for “ ________ Ultra ___________”. Violet This can damage the _________ retina in your eyes, and cause

sunburn and even _______ skin cancer. Its uses include detecting forged ______ bank _______. notes

7. X-rays are used to see inside people, and for _________ airport security.

8. Gamma rays are given off by some ________________ radioactive substances. We can use them to kill ________ cancer cells,

which is called R_______________ adiotherapy .

9. My Quiz score is ____%. 100


10. Name ________________________________

Go to the web site www.darvill.clara.net/emag

Name How they’re made Uses Dangers

Gamma rays

X-rays

Ultra-violet

Visible Light

Infra-red

Microwaves

Radio Waves

Stars radioactive substances

Stars

X-ray machines

Nebula

Sun

Special lamps, sun beds

anything hot enough to glow

light bulbs

Hot objects body

Stars, lamps, flames

Magnetrons "chips"

Extremely high frequency radio waves

Stars, Sparks, and lightning

Transmiters

_____ Frequency _____ frequency,

Short wavelength ______ Wavelength


Learning Goal for this section:

Explain and compare nuclea reactions (radioactive decay, fission and fusion), the energy changes associated

with them and their associated saftey issues.

Notes Section:

6 protons 8 neutrons 14 mass Carbon -> Nitrogen 14 mass 7 protons

Beta Particle

Negative Electron

Positive Electron Positron

Gamma Energy

very high frequency so it can penetrate deeply.

Half-Life 100g of a substance, if it has a 20min half-life, after 20min it has

50g, after another 20min, it has 25g later. The initial substance is conver

into a different element. So that after the first half-life the substance has 5

intitail substance and the other 50g is changed into a different substance


The Nucleus

A typical model of the atom is called the Bohr Model, in

honor of Niels Bohr who proposed the structure in 1913. The Bohr atom consists of a central nucleus

composed of neutrons and protons, which is surrounded by electrons which “orbit” around the nucleus.

Protons carry a positive charge of one and have a mass of about 1 atomic mass unit or amu (1 amu =1.7x10-

27 kg, a very, very small number). Neutrons are electrically “neutral” and also have a mass of about 1 amu. In

contrast electron carry a negative charge and have mass of only 0.00055 amu. The number of protons in a

nucleus determines the element of the atom. For example, the number of protons in uranium is 92 and the

number in neon is 10. The proton number is often referred to as Z.

Atoms with different numbers of protons are called elements, and are arranged in the periodic table with

increasing Z.

Atoms in nature are electrically neutral so the number of electrons orbiting the nucleus equals the number of

protons in the nucleus.

Neutrons make up the remaining mass of the nucleus and provide a means to “glue” the protons in place.

Without neutrons, the nucleus would split apart because the positive protons would repel each other. Elements

can have nucleii with different numbers of neutrons in them. For example hydrogen, which normally only has

one proton in the nucleus, can have a neutron added to its nucleus to from deuterium, ir have two neutrons

added to create tritium, which is radioactive. Atoms of the same element which vary in neutron number are

called isotopes. Some elements have many stable isotopes (tin has 10) while others have only one or two. We

express isotopes with the nomenclature Neon-20 or 20 Ne 10, with twenty representing the total number of

neutrons and protons in the atom, often referred to as A, and 10 representing the number of protons (Z).

Alpha Particle

Decay

Alpha decay is a radioactive process in which a

particle with two neutrons and two protons is

ejected from the nucleus of a radioactive atom. The particle is identical to the nucleus of a helium atom.

Alpha decay only occurs in very heavy elements such as uranium, thorium and radium. The nuclei of these

atoms are very “neutron rich” (i.e. have a lot more neutrons in their nucleus than they do protons) which makes

emission of the alpha particle possible.

After an atom ejects an alpha particle, a new parent atom is formed which has two less neutrons and two less

protons. Thus, when uranium-238 (which has a Z of 92) decays by alpha emission, thorium-234 is created

(which has a Z of 90).

Because alpha particles contain two protons, they have a positive charge of two. Further, alpha particles are

very heavy and very energetic compared to other common types of radiation. These characteristics allow alpha

particles to interact readily with materials they encounter, including air, causing many ionizations in a very short

distance. Typical alpha particles will travel no more than a few centimeters in air and are stopped by a sheet of

paper.


Beta Particle Decay

Beta decay is a radioactive process in which an electron is emitted from the nucleus of a radioactive

atom Because this electron is from the nucleus of the atom, it is called a beta particle to distinguish it

from the electrons which orbit the atom.

Like alpha decay, beta decay occurs in isotopes which are “neutron rich” (i.e. have a lot more

neutrons in their nucleus than they do protons). Atoms which undergo beta decay are located below

the line of stable elements on the chart of the nuclides, and are typically produced in nuclear reactors.

When a nucleus ejects a beta particle, one of the neutrons in the nucleus is transformed into a proton.

Since the number of protons in the nucleus has changed, a new daughter atom is formed which has

one less neutron but one more proton than the parent. For example, when rhenium-187 decays

(which has a Z of 75) by beta decay, osmium-187 is created (which has a Z of 76). Beta particles

have a single negative charge and weigh only a small fraction of a neutron or proton. As a result, beta

particles interact less readily with material than alpha particles. Depending on the beta particles

energy (which depends on the radioactive atom), beta particles will travel up to several meters in air,

and are stopped by thin layers of metal or plastic.

Positron emission or beta plus decay (β+ decay) is a subtype of radioactive decay called beta decay,

in which a proton inside a radionuclide nucleus is converted into a neutron while releasing a positron

and an electron neutrino (νe). Positron emission is mediated by the weak force.

An example of positron emission (β+ decay) is shown with magnesium-23 decaying into sodium-23:

23 Mg12 → 23 Na11 + e +

Because positron emission decreases proton number relative to neutron number, positron decay

happens typically in large "proton-rich" radionuclides. Positron decay results in nuclear transmutation,

changing an atom of one chemical element into an atom of an element with an atomic number that is

less by one unit.

Positron emission should not be confused with electron emission or beta minus decay (β− decay),

which occurs when a neutron turns into a proton and the nucleus emits an electron and an

antineutrino.


Gamma

Radiation

After a decay reaction, the nucleus is often in an

“excited” state. This means that the decay has

resulted in producing a nucleus which still has

excess energy to get rid of. Rather than emitting another beta or alpha particle, this energy is lost by

emitting a pulse of electromagnetic radiation called a gamma ray. The gamma ray is identical in

nature to light or microwaves, but of very high energy.

Like all forms of electromagnetic radiation, the gamma ray has no mass and no charge. Gamma rays

interact with material by colliding with the electrons in the shells of atoms. They lose their energy

slowly in material, being able to travel significant distances before stopping. Depending on their initial

energy, gamma rays can travel from 1 to hundreds of meters in air and can easily go right through

people.

It is important to note that most alpha and beta emitters also emit gamma rays as part of their decay

process. However, their is no such thing as a “pure” gamma emitter. Important gamma emitters

including technetium-99m which is used in nuclear medicine, and cesium-137 which is used for

calibration of nuclear instruments.

Half Life

Half-life is the time required for the quantity of a

radioactive material to be reduced to one-half its

original value.

All radionuclides have a particular half-life, some

of which a very long, while other are extremely

short. For example, uranium-238 has such a

long half life, 4.5x109 years, that only a small fraction has decayed since the earth was formed. In

contrast, carbon-11 has a half-life of only 20 minutes. Since this nuclide has medical applications, it

has to be created where it is being used so that enough will be present to conduct medical studies.


The Learning Goal for this assignment is:

Distinguish between bonding forces holding compound

together and other attractive forces, including hydrogen bonding

and van der waald forces.

Introduction to Ionic Compounds

Those molecules that consist of charged ions with opposite charges are called IONIC. These ionic

compounds are generally solids with high melting points and conduct electrical current. Ionic

compounds are generally formed from metal and a non-metal elements. See Ionic Bonding below.

Ionic Compound Example

For example, you are familiar with the fairly benign unspectacular behavior of common white

crystalline table salt (NaCl). Salt consists of positive sodium ions (Na + ) & negative chloride ions (Cl - ).

On the other hand the element sodium is a silvery gray metal composed of neutral atoms which react

vigorously with water or air. Chlorine as an element is a neutral greenish-yellow, poisonous, diatomic

gas (Cl2).

The main principle to remember is that ions are completely different in physical and chemical

properties from the neutral atoms of the elements.

The notation of the + and - charges on ions is very important as it conveys a definite meaning.

Whereas elements are neutral in charge, IONS have either a positive or negative charge depending

upon whether there is an excess of protons (positive ion) or excess of electrons (negative ion).

Formation of Positive Ions

Metals usually have 1-4 electrons in the outer energy level. The electron arrangement of a rare gas is

most easily achieved by losing the few electrons in the newly started energy level. The number of

electrons lost must bring the electron number "down to" that of a prior rare gas.

How will sodium complete its octet?

First examine the electron arrangement of the atom. The atomic number is eleven, therefore, there

are eleven electrons and eleven protons on the neutral sodium atom. Here is the Bohr diagram and

Lewis symbol for sodium:


This analysis shows that sodium has only one electron in its outer level. The nearest rare gas is neon

with 8 electron in the outer energy level. Therefore, this electron is lost so that there are now eight

electrons in the outer energy level, and the Bohr diagrams and Lewis symbols for sodium ion and

neon are identical. The octet rule is satisfied.

Ion Charge?

What is the charge on sodium ion as a result of losing one electron? A comparison of the atom and

the ion will yield this answer.

Sodium Atom

Sodium Ion

11 p+ to revert to 11 p + Protons are identical in

12 n an octet 12 n

the atom and ion.

Positive charge is

11 e- lose 1 electron 10 e-

caused by lack of

0 charge + 1 charge

electrons.

Formation of Negative Ions

How will fluorine complete its octet?

First examine the electron arrangement of the atom. The atomic number is nine, therefore, there are

nine electrons and nine protons on the neutral fluorine atom. Here is the Bohr diagram and Lewis

symbol for fluorine:

This analysis shows that fluorine already has seven electrons in its outer level. The nearest rare gas

is neon with 8 electron in the outer energy level. Therefore only one additional electron is needed to

complete the octet in the fluorine atom to make the fluoride ion. If the one electron is added, the Bohr

diagrams and Lewis symbols for fluorine and neon are identical. The octet rule is satisfied.


Ion Charge?

What is the charge on fluorine as a result of adding one electron? A comparison of the atom and the

ion will yield this answer.

Fluorine Atom Fluoride Ion *

9 p+ to complete 9 p + Protons are identical in

10 n octet 10 n

9 e- add 1 electron 10 e-

0 charge - 1 charge

the atom and ion.

Negative charge is

caused by excess

electrons

* The "ide" ending in the name signifies a simple negative ion.

Summary Principle of Ionic Compounds

An ionic compound is formed by the complete transfer of electrons from a metal to a nonmetal and

the resulting ions have achieved an octet. The protons do not change. Metal atoms in Groups 1-3

lose electrons to non-metal atoms with 5-7 electrons missing in the outer level. Non-metals gain 1-4

electrons to complete an octet.

Octet Rule

Elemental atoms generally lose, gain, or share electrons with other atoms in order to achieve the

same electron structure as the nearest rare gas with eight electrons in the outer level.

The proper application of the Octet Rule provides valuable assistance in predicting and explaining

various aspects of chemical formulas.

Introduction to Ionic Bonding

Ionic bonding is best treated using a simple

electrostatic model. The electrostatic model

is simply an application of the charge

principles that opposite charges attract and

similar charges repel. An ionic compound

results from the interaction of a positive and

negative ion, such as sodium and chloride in

common salt.

The IONIC BOND results as a balance

between the force of attraction between

opposite plus and minus charges of the ions

and the force of repulsion between similar

negative charges in the electron clouds. In

crystalline compounds this net balance of

forces is called the LATTICE ENERGY.

Lattice energy is the energy released in the

formation of an ionic compound.

DEFINITION: The formation of an IONIC

BOND is the result of the transfer of one or

more electrons from a metal onto a nonmetal.


Metals, with only a few electrons in the outer energy level, tend to lose electrons most readily. The

energy required to remove an electron from a neutral atom is called the IONIZATION POTENTIAL.

Energy + Metal Atom ---> Metal (+) ion + e-

Non-metals, which lack only one or two electrons in the outer energy level have little tendency to lose

electrons - the ionization potential would be very high. Instead non-metals have a tendency to gain

electrons. The ELECTRON AFFINITY is the energy given off by an atom when it gains electrons.

Non-metal Atom + e- --- Non-metal (-) ion + energy

The energy required to produce positive ions (ionization potential) is roughly balanced by the energy

given off to produce negative ions (electron affinity). The energy released by the net force of

attraction by the ions provides the overall stabilizing energy of the compound.

Notes Section:


The Learning Goal for this assignment is:

Introduction to Covalent Bonding:

Bonding between non-metals consists of two electrons shared between two atoms. Using the Wave

Theory, the covalent bond involves an overlap of the electron clouds from each atom. The electrons

are concentrated in the region between the two atoms. In covalent bonding, the two electrons shared

by the atoms are attracted to the nucleus of both atoms. Neither atom completely loses or gains

electrons as in ionic bonding.

There are two types of covalent bonding:

1. Non-polar bonding with an equal sharing of electrons.

2. Polar bonding with an unequal sharing of electrons. The number of shared electrons depends on

the number of electrons needed to complete the octet.

NON-POLAR BONDING results when two identical non-metals equally share electrons between

them. One well known exception to the identical atom rule is the combination of carbon and hydrogen

in all organic compounds.

Hydrogen

The simplest non-polar covalent molecule is hydrogen. Each hydrogen

atom has one electron and needs two to complete its first energy level.

Since both hydrogen atoms are identical, neither atom will be able to

dominate in the control of the electrons. The electrons are therefore

shared equally. The hydrogen covalent bond can be represented in a

variety of ways as shown here:

The "octet" for hydrogen is only 2 electrons since the nearest rare gas is

He. The diatomic molecule is formed because individual hydrogen atoms

containing only a single electron are unstable. Since both atoms are

identical a complete transfer of electrons as in ionic bonding is

impossible.

Instead the two hydrogen atoms SHARE both electrons equally.

Oxygen

Molecules of oxygen, present in about 20% concentration in air are

also covalent molecules. See the graphic on the left of the Lewis Dot

Structure.

There are 6 electrons in the outer shell, therefore, 2 electrons are

needed to complete the octet. The two oxygen atoms share a total of

four electrons in two separate bonds, called double bonds.

The two oxygen atoms equally share the four electrons.


POLAR BONDING results when two different non-metals unequally share electrons between them.

One well known exception to the identical atom rule is the combination of carbon and hydrogen in all

organic compounds.

The non-metal closer to fluorine in the Periodic Table has a greater tendency to keep its own electron

and also draw away the other atom's electron. It is NOT completely successful. As a result, only

partial charges are established. One atom becomes partially positive since it has lost control of its

electron some of the time. The other atom becomes partially negative since it gains electron some of

the time.

Hydrogen Chloride

Hydrogen Chloride forms a polar covalent molecule. The graphic

on the left shows that chlorine has 7 electrons in the outer shell.

Hydrogen has one electron in its outer energy shell. Since 8

electrons are needed for an octet, they share the electrons.

However, chlorine gets an unequal share of the two electrons,

although the electrons are still shared (not transferred as in ionic

bonding), the sharing is unequal. The electrons spends more of the

time closer to chlorine. As a result, the chlorine acquires a "partial"

negative charge. At the same time, since hydrogen loses the

electron most - but not all of the time, it acquires a "partial" charge.

The partial charge is denoted with a small Greek symbol for delta.

Water

Water, the most universal compound on all of the earth, has the property of

being a polar molecule. As a result of this property, the physical and

chemical properties of the compound are fairly unique.

Dihydrogen Oxide or water forms a polar covalent molecule. The graphic on

the left shows that oxygen has 6 electrons in the outer shell. Hydrogen has

one electron in its outer energy shell. Since 8 electrons are needed for an

octet, they share the electrons.

Notes Section:

1. Count the Valence e-

2.Find central atom and bond other atom to it. Submit bonds from total and put lone pairs and double bonds and

triple bonds as needed.

3. Find formal charges. Try tp get as close to zero as possible.


C 2 H 6 O Ethanol CH 3 CH 2 O

Step 1

Find valence e- for all atoms. Add them together.

C: 4 x 2 = 8

H: 1 x 6 = 6

O: 6

Total = 20

Step 2

Find octet e- for each atom and add them together.

C: 8 x 2 = 16

H: 2 x 6 = 12

O: 8

Total = 36

Step 3

Subtract Step 1 total from Step 2.

Gives you bonding e-.

36 – 20 = 16e-

Step 4

Find number of bonds by diving the number in step 3 by 2

(because each bond is made of 2 e-)

16e- / 2 = 8 bond pairs

These can be single, double or triple bonds.

Step 5

Determine which is the central atom

Find the one that is the least electronegative.

Use the periodic table and find the one farthest

away from Fluorine or

The one that only has 1 atom.


Step 6

Put the atoms in the structure that you think it will

have and bond them together.

Put Single bonds between atoms.

Step 7

Find the number of nonbonding (lone pairs) e-.

Subtract step 3 number from step 1.

20 – 16 = 4e- = 2 lone pairs

Step 8

Complete the Octet Rule by adding the lone

pairs.

Then, if needed, use any lone pairs to make

double and triple bonds so that all atoms meet

the Octet Rule.

See Step 4 for total number of bonds.


Linear

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp AX 2 None 180

BeCl 2

Beryllium Dichloride

CI

Be

CI

element bond lone pair

C


Trigonal Planar

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 2 AX 3 None 120

BF 3

Boron Trifluoride

F

B

F

F

element bond lone pair

C


Bent

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 AX 2 E 2 2 104.5

OF 2

Oxygen Difluoride

O

F

F

element bond lone pair

C


Tetrahedral

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 AX 2 None 109.5

Phosphate

PO 4

3-

O

O

P

O

O

element bond lone pair

C


Bent

Molecular Geometry

Orbital Equation Lone Pairs Angle

Sp 3 AX 2 E 2 1 107

Phosphorus

Trihydride

H

PH 3

P H

C

element bond lone pair

C


Bent

Molecular Geometry

Orbital Equation Lone Pairs Angle

Sp 2 AX 2 E 1 116

Trioxide

O 3

O O O

element bond lone pair

C


Trigonal Bi Pyramidal

Molecular Geometry

Orbital Equation Lone Pairs Angle

Sp 3 d AX 5 None 120/90

Phosphorus

pentachloride

PCl 5

Cl

Cl

Cl

P

Cl

Cl

element bond lone pair

C


T-Shaped

Molecular Geometry

Orbital Equation Lone Pairs Angle

Sp 3 d Ax 3 E 2 2 90

Chlorine

Trifluoride

ClF 3

F

Cl

F

F

element bond lone pair

C


Octahedral

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 d 2 AX 6 None 90

Sulfur hexafluoride

SF 6

F

F

F

S

F

F

F

element bond lone pair

C


Square Planar

Molecular Geometry

Orbital Equation Lone Pairs Angle

sp 3 d 2 AX 4 E 2 2 90

Iodine Tetrachloride ion

ICl 4

-

C

C

I

C

C

element bond lone pair

C


Orbitals Equation Lone Pairs Angle

Name

sp AX2 None 180

Linear

sp 2 AX3 None 120

Trigonal Planar

sp 2 AX2E 1 116

Bent

sp 3 AX4 None 109.5

Tetrahedral

sp 3 AX3E 1 107

Trig. Pyramidal

sp 3 AX2E2 2 104.5

Bent

sp 3 d AX5 None 120/90

Trig. Bipyramidal

sp 3 d AX3E2 2 90

T-Shaped

sp 3 d 2 AX6 None 90

Octahedral

sp 3 d AX4E2 1 90

Square Planar


Name Formula Charge

Dichromate Cr₂O₇ 2-

Sulfate SO₄ 2-

Hydrogen Carbonate HCO₃ 1-

Hypochlorite ClO 1-

Phosphate PO₄ 3-

Nitrite NO₂ 1-

Chlorite ClO₂ 1-

Dihydrogen phosphate H₂PO₄ 1-

Chromate CrO₄ 2-

Carbonate CO₃ 2-

Hydroxide OH 1-

Hydrogen phosphate HPO₄ 2-

Ammonium NH₄ 1+

Acetate C₂H₃O₂ 1-

Perchlorate ClO₄ 1-

Permanganate MnO₄ 1-

Chlorate ClO₃ 1-

Hydrogen Sulfate HSO₄ 1-

Phosphite PO₃ 3-

Sulfite SO₃ 2-

Silicate SiO₃ 2-

Nitrate NO₃ 1-

Hydrogen Sulfite HSO₃ 1-

Oxalate C₂O₄ 2-

Cyanide CN 1-

Hydronium H₃O 1+

Thiosulfate S₂O₃ 2-


Chapter 9

Unit 4

Chemical Names and Formulas

The students will learn how the periodic table helps them

determine the names and formulas of ions and compounds.

Chapter 22 Hydrocarbon Compounds

The student will learn how Hydrocarbons are named and the

general properties of Hydrocarbons.

Describe how different natural resources are produced and how their rates

of use and renewal limit availability.

Students will explore local, national, and global renewable and nonrenewable

resources.

Students will explain the environmental costs of the use of renewable and

nonrenewable resources.

Students will explain the benefits of renewable and nonrenewable resources.

Nuclear reactors

Natural gas

Petroleum

Refining

Coal


Chapter 23 Functional Groups

The student will learn what effects functional groups have on

organic compounds and how chemical reactions are used in

organic compounds.

Describe the properties of the carbon atom that make the diversity of carbon

compounds possible.

Identify selected functional groups and relate how they contribute to

properties of carbon compounds.

Students will identify examples of important carbon based molecules.

Students will create 2D or 3D models of carbon molecules and explain why this

molecule is important to life.

covalent bond

single bond

double bond

triple bond

monomer

polymer


http://www.bbc.co.uk/education/guides/zm9hvcw/revision

methane- is a natural gas used for cooking and heating

propane- is gas used in gas cylinders for BBQ etc

octane- used in petrol for cars

Parent Molecule- the longest unbranched

chain containing the functional group.

The position of the functional group is labled

with a number

Methane - Monsters

Ethane - Eat The rule is a comma between numbers, and a

Propane - Pupils dash between # and letters.

Butane - But

Pentane - Prefer

Hexane - Hairy

Alkenes- all end in -ene

Heptane - Haggis

all contain a carbon to carbon

Octane - Occasionally

double bond-unsaturated

only contain single bonds

saturated.

uses are fuels, solvents,

plastics, and vinegar

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