Homework

• Read Textbook Pgs.291 – 303– Matching – p. 304– Review p. 304 #’s 2-12

Atoms, Elements, and thePeriodic Table

Chapters 17-19

The Chemistry of Life

• Cell (skin cell)• basic building block of

living organisms (life)

• Atom • basic building block of

matter

• Where can we find a list of all the different types of atoms?

The Periodic Table of Elements

What are all these letters?

Atomic Symbols

What do all these letters represent?

Atoms? Elements? Both?

How do we distinguish between an element and an atom?

The Periodic Table of Elements

Atom is used to refer to submicroscopic particles is a sample and element is used for

microscopic and macroscopic samples.

Label Each Structure

1._________

1.Electron

2. ___________

2. Proton

3. __________

3. Neutron

• The parts that make up an atom are called subatomic subatomic particlesparticles.– Protons (p+) positively charged particle– Neutron (n) neutral particle (uncharged)– Electrons (e-) negatively charged particle

• Neutrons and Protons are located in the nucleusnucleus of an atom and are called nucleons.

• Electrons orbit around the nucleus.

• How are atoms of different elements distinguished from one another? In other words, how do we distinguish a helium atom from a carbon atom?– Their Number of Protons.

• What on the periodic table give us this information?– Atomic Number

Carbon

Atomic Number = 6

Mass Number = 12

The Periodic Table and Atoms

• Atomic Number– Gives us the number

of protons. and electrons in a neutral atom.

• Atom Mass– Tells us the atomic

mass of an atom of that element (measured in AMU, we’ll talk about this later)

How do we know how many neutrons there are in an atom?

• 1 proton = 1 atomic mass unit

• 1 neutron = 1 atomic mass unit

• 1 electron = nothing– Electrons are 1/1840 the

size of a proton or neutron.

– They do not factor into the atomic mass of an element

•So what formula can we come up with to determine the number of neutrons in an atom?

•Atomic Mass – Atomic Number = # of Neutrons

Lets practice!!! Find the missing information?

Element Atomic

#

Atomic

Mass

Protons Electons Neutrons

Ar 18 39.948 amu

He 2 2

O 15.999 amu

8

IsotopesOrganization of the Periodic Table

andProperties of the Elements

• Homework– Review questions 13-20 page 304.

• Do Now– A cat strolls across your backyard. An hour later, a dog

with its nose to the ground follows the trail of the cat. Explain what is going on in terms of atoms.

– Which are older, the atoms in the body of an elderly person or those in the body of a baby?

Isotopes

•    Because they have the same number of protons, all isotopes of an element have the same chemical properties.

Radioactive Isotopes

• Radioactive decay – the break down of the nucleus of an atom

• Radioactive Isotopes– Isotopes where the nucleus is unstable and

breaks down at a constant rate over time– Radioactive isotopes can be useful is a variety

of ways

Homework

• Study for quiz on Chapter 17– Multiple Choice – Bring Pencil

– If an atom has 43 electrons, 56 neutrons, and 43 protons, what is its approximate atomic mass? What is the element?

– Evidence of the existence of neutrons did not come until many years after the discovery of electrons and protons. Give a possible explaination.

– Which has more atoms: a one gram sample of carbon 12 or a one gram sample of carbon 13? Explain.

– Shiny and Opaque– Good Conductors of Electricity and Heat– Malleable and Ductile– Most solid at room temperature

• Exceptions – mercury, gallium, cesium, francium

– Hydrogen – “A little different”• Typically a nonmetal but at high pressures hydrogen

solidifies and displays metallic properties• Under normal condition hydrogen is a nonmetallic

GAS!!!

Metals

Nonmetals

– Poor Conductors of ______ and ______.– May be transparent– Brittle– At room Temperature

• Some are solids (Carbon)

• Some are liquids (Bromine)

• Some are gases (Helium)

Metalloids

• Only six elements considered metalloids– Boron (B), Silicon (Si), Germanium (Ge),

Arsenic (As), Tin (Sn), Antimony (Sb)– Have both metallic and nonmetallic properties– Weak conductors

• Used as “semiconductors”

• Period – horizontal row – properties gradually change as you move from element

to element • Periodic Trend

• Groups – vertical column– Properties of elements in the same group are similar– There are 6 main vertical groups

• Alkali Metals• Alkaline-earth Metals• Transitional Metals• Chalcogens• Halogens• Noble gasses

• Alkali Metals – Group I- form alkaline solutions - (al-qali – ashes) – ashes of alkali metals (with water) were used as hand cleaners

• Alkaline-earth Metals – Group II - form alkaline solutions – some elements have high melting points and are not affected by fire

• Chalcogens – (ore-forming) – these elements are commonly found in ores

• Halogens – (salt-forming) – these elements commonly form salts

• Noble Gases – unreactive gases “Noble” don’t mix with the “common folk”

• Transition Metals – metals which do not form alkaline solutions with water – these metals tend to be harder and less reactive with water than the alkali metals– Lanthanides – follow (La) Lanthanum– Actinides – follow (Ac) Actinum

Homework

• Read p. 307-318– Matching p. 319– Review p. 319 #’s 1-9

Atomic Models

• Distinguish between a PHYSICAL and CONCEPTUAL model

• Explain why scientists know the identity of the chemical composition of stars

• Explain how the quantum hypothesis explains atomic spectra

• Explain the usefulness of the shell model of the atom

Models

• Physical model – representation of an object on a different scale

• Conceptual model – representation of a system that helps us predict how the system behaves

Neils Bohr and the

Quantum Hypothesis• Quantum Hypothesis –

states that a beam of light is not a continuous train of waves, but is a stream of zillions of discrete particles (photons)

• Quantum – the smallest unit of something

Electrons and Potential Energy

Bohr and the Principal Quantum Number

• Principal Quantum Number - An integer that specifies the quantized energy level of an atomic orbital

• Planetary Model

Atoms and Light

• Atomic Spectra – the pattern of frequencies (light) formed by a given element

Color Element [common compounds]

  Red  [Pictures]

      Lithium [Li2CO3],

Strontium [SrCO3, Sr(NO3)2]

  Orange  [Pictures]

      Calcium [CaCO3, CaSO4]

  Yellow  [Pictures]

      Sodium [NaCl, NaNO3,

Na2CO3]

  Yellowish-Green

      Boron [Borax - Na2B4O7]

  Green  [Pictures]

      Barium [Ba(NO3)2, BaCl2],

Copper [CuSO4]

  Blue  [Pictures]

      Copper [CuO, CuCO3], Copper Halides  [CuCl2]

  Purple-Violet  [Pictures]

      Potassium [KClO3, KCl,

KNO3, K2SO4]

  White-Silver

      Aluminum, Magnesium, Titanium

Hg atomic spectra (above)

Ne atomic spectra (Left)

Shell Model and

arrangement of e-

• Shell – a region of space in which an electron may be located around the nucleus

• Valence electrons – electrons on the outermost orbital

Homework Review Questions

9-16

Draw Your Own Bohr-Model of these Atoms

• We simplify the orbitals when we draw a Bohr Model.• We just use rings!!• Orbital 1 – 2e• Orbital 2 – 8e• Orbital 3 – 8e• Orbital 4 – 18e• Orbital 5 – 18e• Orbital 6 – 32e• Orbital 7 – 32e• Total 118 Electrons

Radioactivity and Imbalance of Forces

• Strong nuclear force - The Strong Nuclear Force (also referred to as the strong force) is one of the four basic forces in nature (the others being gravity, the electromagnetic force, and the weak nuclear force). As its name implies, it is the strongest of the four. However, it also has the shortest range, meaning that particles must be extremely close before its effects are felt. Its main job is to hold together the subatomic particles of the nucleus (protons, which carry a positive charge, and neutrons, which carry no charge. These particles are collectively called nucleons).

• Balance between electric and the strong nuclear force is needed for a stable nucleus.

Radioactivity

Radioactivity – the process by which certain elements emit particular form of radiation.

• Radioactive – any element which emits any of these forms of radiation– Alpha radiation– Beta radiation– Gamma radiation

Alpha Radiation• Alpha particle –

combination of two protons and two neutrons

Beta Radiation• Beta particle – an

electron that is ejected by an atom

– Beta-minus decay• Neutron -> proton

emitting electron and antineutrino

– Beta-plus decay• Proton -> neutron

emitting neutrino and positron

Transmutation• The changing of

one element to another

Radioactive Decay (Half-Life)

• Homework– P. 368 Matching

– Test Tuesday

– Review Class Monday

Today we will learn about:

1. Properties of Matter• Physical Properties

• Chemical Properties

2. Physical Changes

3. Chemical Reactions (chemical changes)

H + O => H2O

Chapter 21: Chemistry

Elements of ChemistryChapter 21

• What is chemistry?– Chemistry is the study of matter and the

transformations it can undergo.

– In other words, Chemistry is the study of atoms, molecules, and the interactions between them.

• Molecule – a unit of matter consisting of 2 or more atoms linked together

Properties of Substances Physical and Chemical

• Physical properties – describe the look or feel of a substance– Color, density, texture,

phase

• Physical change – substance changes its phase or some other physical property but not its chemical composition

Properties of SubstancesPhysical and Chemical

• Chemical properties – properties that relate to how a substance reacts with others.– Chemical property of baking soda is that it

reacts with vinegar to produce CO2 and H2O.

• Chemical change – any change in a substance that involves the rearrangement of its chemical bonds– Chemical bond – attraction between two

atoms that hold them together in a molecule.

Chemical Reaction

• Chemical reaction – same thing as chemical change.

• Let’s take a look at the chemical reaction that magnesium (Mg) undergoes with oxygen (O).

Magnesium

• Draw the electron-dot structure of Mg

Oxygen• Draw the

electron-dot structure of Oxygen.

Chapter 23

• Homework – Page 368Matching

– Review 6-15

• Homework for tomorrow night – Review 16-21

• After today we will be able to distinguishing:– between a chemical

change and a physical change.

– Explain the difference between an element and a compound

– Create and balance chemical formula

Chemical Change or Physical Change

• Physical change – substance changes its phase or some other physical property but not its chemical composition

• Chemical change – any change in a substance that involves the rearrangement of its chemical bonds

Element or Compound

How will Magnesium (Mg) and Oxygen (O) bind together?

Let’s take a look at this reaction.

Magnesium Oxide (MgO)

Can we come up with an easier way to represent this

Chemical Reaction?• Yes, we can write a chemical equation.

– chemical equation – a representation of a chemical reaction showing the relative numbers of reactants and products

• Mg2+ + O2- MgO

• So we get

• Mg + O MgO

Chemical Equation

• Lets write the chemical equation for Beryllium (Be) reacting with Fluorine (F).

– Be2+ + F1- BeF2

– Be + F BeF2

– Can we create or destroy matter?

• No, So we need to balance our equation then.

– Be + 2F BeF2

Balance the Following Chemical Equations

• Hydrogen reacts with oxygen to form water

• ___H2 + ___O2 => ___H2O

• ___Na + ___O2 => ___Na2O2

• ___Mg + ___Cl2 => ___MgCl2

• Write the complete and balanced chemical reaction for Sodium reacting with Sulfur.

Chemical Equations used to describe

Chemical Reactions

Homework

Media Workbook

Pages 46 + 47

Chemical Reactions and Equations

• ReactantsReactants - The elements or compounds that enter into a chemical reaction

• ProductsProducts - The elements or compounds produced by a chemical reaction

The Octet Rule• The octet rule says that atoms tend to gain,

lose or share electrons so as to have eight electrons in their outer electron shell. It is a very useful rule but you should also know that there are many bonding situations where it does not apply.

• Very simply, when atoms bond together they want to have 8 electrons in their outer shell.

Chemical Equations

• Step one – Write atomic symbols for Reactants• Step two – Write in the ionic charges for the

reactants• Step three – Write in the product (remember to

cross the ionic charges)• Step four – Balance the equation• Step five – Re-write the complete chemical

equation

MixturesChapter 22

Homework Handout

Mixtures

• A mixture is a combination of two or more substances in which each substance retains its properties.– Pure Material – (not a mixture) a substance

which consists of only a single element or compound

– Impure Material – (mixture) – contains two or more elements or compounds.

Mixtures• Mixtures can

be separated by physical means.

Homework

Classifying Matter

Concentrations

Moles????

Types of Mixtures

• Heterogeneous Mixture – the different components can be seen as individual substances. (Different components can be seen with the naked eye)

• Homogeneous Mixture – Uniform composition throughout. (Different components are mixed at a much finer level and are not readily distinguishable) – Solution – Homogeneous mixture with all components in

the same phase– Suspension – Homogeneous mixture when different

components are in different phases. – Now, Heterogeneous, Homogeneous, Solution, Suspension?

Solutions• Solvent – the component with the largest amount. The component which does the dissolving

• Solute – the component which is dissolved– Saturated solution – a solution

where no more solute can be dissolved

– Unsaturated solution – a solution that has not reached the limit of solute

– Concentration - the amount of solute dissolved per amount of solution.

• Concentration = solute (g)/ solvent (mL)

How much salt can 50mL of water hold???

• Task 1– How much salt can 50mL of water hold?– Solve for the concentration of your solution.

• Question from before. – What did you create when you mixed alcohol

and water?– Why did the volumes not add up?

Moles

• Homework – Handout

• Today you will learn about:– Moles– Mole conversions

• Find the concentration of each solution.

– 23 grams of sugar dissolved in 130mL of water

– 0.12 kg dissolved in 5500 mL of water

• Which will have more particles?– One gram of water

– One gram of glucose

The Mole

• Chemists are often more interested in the number of particles in a solution rather than the number of grams.

• The problem with this is that the number of particles in a solution will be a GINORMOUS number

• This is where the mole comes in.

Mole• One mole of any type of particle is, by definition

6.02 x 1023 particles. (Avagadro’s number)• Similar ideas

– A couple of knuckleheads = 2 Maxes

– A dozen of donuts = 12 donuts

– A mole of baseballs = 6.02 x 1023 Baseballs

– A mole of cars = 6.02 x 1023 Cars

– How many moles of wheels will there be?

– How many wheels will there be?

– A mole of water molecules = how many molecules?

– How many moles of Hydrogen will there be?

– How many particles of Hydrogen will there be?

Moles to Particles

• How many molecules are there in 2.5moles of methane (CH4)?

• How many particles of methane are there in 3.0 moles of methane (CH4)?

• How many atoms of hydrogen are in 3.0 moles of Methane (CH4)?

• How many total atom are there in 3.0 moles (CH4)

Particles to moles

• You have 3.5 x 1023 molecules of water. How many moles do you have?

• You have 6.9 x 1024 atoms of U. How many moles do you have?

Moles and MolarityIPC 1 – Quiz Tuesday Chapter 22

IPC 2 & 3 – Quiz Wednesday Chapter 22

Homework (In Notebook)How many moles of glucose is there in 2.5L of a 3

molar solution? How many molecules of glucose is this?

Agenda1. Go over quiz2. Review Moles3. Learn about Molarity4. Problems on Molarity

• Do Now (Moles)

1. How many atoms of Hg are there in 1.24 moles of Hg?

2. How many moles of C do you have if there are 1.23x10^24 atoms of carbon?

3. How many atoms of oxygen are there in 3moles of carbon dioxide (CO2)?

Molarity

• Concentration = grams solute / liters of solution

• Molarity is concentration using moles

• Molarity = # of Moles solute / liters of solution

Molarity

• What is the molarity of a 2.4L solution containing 1.2 moles of sugar?

• What is the molarity of a 4.0L solution containing 1.2 x 10^24 molecules of HCL?

• How many moles of salt are there in 3L of a 3 molar solution of salt water? How many formula units of salt are there?

Chapter 23(Last day of New Material)

• TEST– Chapter 21, 22, 23

– IPC 1 & 2 – Test Friday

– IPC 3 Test Thursday

– Homework • Molar Mass Handout

• Agenda– Go over moles and

molar mass handout

– New Material• Ionic Bonds

• Covalent Bonds

• Molar Mass

– Molar Mass Handout

Ionic Bonds

• IONIC Bonds - An electric force of attraction between two oppositely charged ions.

• Ionic compound – all compounds containing ions.

Covalent Bonds

• Covalent Bond – Type of electrical attraction in which atoms are held together by their mutual attraction for shared electrons.

• Covalent Compound – a substance that is made up of atoms held together by covalent bonds. A substance made of molecules (as opposed to ions) is a covalent compound.

Molar Mass

• Molar Mass – the mass of one mole of a substance.

• What is the molar mass of – Bromine (Br)

– Tin (Sn)

– Aluminum Oxide (Al2O3)

– Carbon Dioxide (CO2)

Task

• Get me one mole of water?????

Homework

• Create and solve your own mole problem!!!

Molar Mass

• What is the mass (g) of 2.37Moles of water?

• You have 440g of methane. – How many moles of methane do you have?

• How many molecules of methane do you have?

• How many atoms of hydrogen do you have?

Isotopes• The properties of an atom are determined by the

number of protons.• If the number of protons in the atom of an element

are changed a new element has been formed.• However, the number of neutrons in the atom of

an element can be change and still have the same element.

• Isotopes - Atoms of the same element with different numbers of neutrons which gives the atoms different masses.

• What about electrons?

If all of these carbon atoms have atomic masses of 12, 13, or 14, then

why is the atomic mass of carbon 12.011amu?

Homework

• Read textbook pgs. 40-48

• Handout on 2-1

Homework

• Read textbook pgs 44-48

• Study Guide 2-2

Quick Quiz

1. How many electron, protons and neutrons are in:

1. Hydrogen – 1

2. Silver - 108

3. Lithium – 7

4. Calcium - 40

2. What is a radioactive isotope? Describe two scientific uses of radioactive isotopes.

• SubstanceSubstance– That which has mass and occupies space; matter.

– A material of a particular kind or constitution.

• A Chemical CompoundChemical Compound is a substance formed by the chemical combination of two r more elements is definite proportion.

• MoleculeMolecule – the smallest form of a compound

Chemical Compounds

Sodium• As a pure substance sodium is a soft metal

which can be easily cut.• Its atomic number is 11.

When sodium is added to water it reacts violently.

Chlorine• As a pure substance chlorine is a poisonous

gas.• Its atomic number is 17.

What do we get if we combine Na and CL.

Ionic Bonds

• A bond between atoms when electrons are transferred.

Covalent Bonds• A bond between atoms when electrons are

shared.

Types of covalent bonds:

-single bond

-double bond

-triple bond

Homework

• Read textbook pgs 44-48

• Study Guide 2-2

Quiz

1. Explain the relationship among atoms, elements, and compounds.

2. What is a radioactive isotope? Describe two scientific uses of radioactive isotopes.

3. How are atoms in a compound held together?

4. Distinguish between ionic bonds and covalent bonds.

WATER (H2O)

Electronegativity and Polarity

• Electronegativity – the affinity (want) for electrons.

•    A water molecule is polar because there is an uneven distribution of electrons between the oxygen and hydrogen atoms.

TIMEOUT from WATER

Van Der Waals Forces

• Forces of attraction between molecules.

• Not nearly as strong as ionic or covalent bonds.

• Molecule – smallest unit of most compounds.

Hydrogen Bonds

• A type of Van Der Waals force.

• Not as strong as Ionic or covalent bonds

• But the strongest of all other bonds.

CohesionThe attraction between molecules of the same

substance

AdhesionThe attraction between molecules of different

substances.

Capillary ActionCombination of adhesion and cohesion

Why does polarity and hydrogen bonding makes water so special?

There are many reasons, but two main reasons are:

#1 ICE Floats

• ICE FLOATS• Water is one of the only substances that

expands when is freezes.• Almost all other substance constricts when

temperature drops.• This occurs because of the polarity of the

molecule.• Ice is less dense than water.

Homework

• Study Section 2-1, 2-2

• 30pt Quiz Tomorrow

#2 - Water is the“UNIVERSAL SOLVENT”

• Water has the ability to dissolve a vast amount of different substances.

• Mixture – material composes of 2 or more elements (or compounds) which can be physically separated. – Two types of mixtures are

• Solutions – mixture where all components are evenly distributed (salt water)

• Suspensions – mixture of water that contains nondissolved materials (pond water)

Solutions• Solute – the substance that is dissolved.

• Solvent – the substance that dissolves the solute.

Suspension

Water (Acids and Bases)What makes up water?

HH (hydrogen) + OO (oxygen)

What is the chemical formula for water?

HH22OO

What is the chemical equation for water?

2H2H + OO HH22OO

Do you think water can do this?

HH22OO 2H2H + OO

Yes, but its more likely to do this:Yes, but its more likely to do this:

HH22O O H H++ + OH + OH--

In pure water, about 1 water molecule in In pure water, about 1 water molecule in 550 million will react to form hydrogen 550 million will react to form hydrogen

ions and hydroxide ions ions and hydroxide ions

The pH Scale

• measurement system used to indicate the concentration of hydrogen ions (H+) in solution; ranges from 0 to 14

Acids   Acidic solutions contain higher concentrations

of H+ ions than pure water and have pH values below 7.

Bases

   Basic, or alkaline, solutions contain lower concentrations of H+ ions than pure water and have pH values above 7.

Acid and Base Buffers

• Cells in our body like to keep their pH between 6.5 and 7.5 on the pH scale.

• A buffer helps to do just that!

• Buffer - weak acid or base that can react with strong acids or bases to help prevent sharp, sudden changes in pH

Homework Read Textbook 49-53Study Guide 2-3 for Wednesday Quiz 2-1 and 2-2 on Wednesday

• Describe the relationship between Acids, Bases, and the pH scale.

• Why is water the universal solvent? How does water dissolve salt?

Organic Chemistry

• The study of carbon compounds.Is Carbon that important to devote an entire branch

of chemistry to it?YESCarbon atoms have four valence electrons. Each

electron can join with an electron from another atom to form a strong covalent bond. Carbon can bond with many elements, including hydrogen, oxygen, phosphorus, sulfur, and nitrogen.

Even more important, carbon can bond with itself.

Macromolecules

Are “giant molecules” formed from thousands or hundreds of thousands of smaller molecules

Monomers (links) are the smaller units which are joined together to form Polymers (bike chain)

The four main groups of macromolecules (organic compounds) are:

Carbohydrates

Lipids

Nucleic Acids

Proteins

Polymerization Polymerization When small molecules, called monomers, join together, they form polymers, or large molecules.

Carbohydrate• Contains C, H, and O usually in a ratio of 1:2:1

• Living things use carbohydrates as their main source of energy. Plants, fungus, and some animals use carbohydrates for structural purposes.

• Monosaccharides(monomer):Monosaccharides(monomer): glucose [C6H12O6], fructose, ribose, deoxyribose

• DisaccharidesDisaccharides: sucrose, lactose- a carb formed by the covalent bonding of 2 or more Monosaccharides

• Polysaccharides(Polymer):Polysaccharides(Polymer): starch, cellulose - large macromolecule formed from monosaccharides includes starches, cellulose and glycogen.

Monosaccharides

Disaccharide

Polysaccharide

Cellulose

Chitin

Homework read pgs 49-53Study Guide 2-3

• Label the following with one or more or the following:

• monomer, polymer, monosaccharide, disaccharide, polysaccharide, simple sugar, starch

• 1. Exoskeleton 2. Cell Wall 3. sugar for coffee

Dehydration Synthesis and Hydrolysis

• Dehydration Synthesis

• Joining the sugars (monomers) to form a disaccharide

Fat• Contains chains of mostly C and H (fatty acids tails) with a

backbone containing C, H, and O (glycerol head)

• Lipids can be used to store energy . Some lipids are important parts of biological membranes and waterproof coverings.

• Lipid, triacylglycerol (circulating fat), triglyceride(storage form of fat), phospholipid(cell membrane), steroid

• Saturated Fat – completely saturated with hydrogen no double bonds

• unsaturated fat – contains one double bond

• polyunsaturated fat – contains more than one double bond

• Fatty Acid Tails are - Hydrophobic

• Glycerol Heads (backbone) are - Hydrophylic

Dehydration Synthesis of Dehydration Synthesis of LipidLipid

Lipids

• Saturated fat – full of hydrogen, no double bonds

• Unsaturated fat – one double bond

• Polyunsaturated fat – more than one double bond

Homework Textbook read 49-53

• 1.  Key Concept Name four groups of organic compounds found in living things.

• 2.  Key Concept Describe at least one function of carbohydrates and lipids.

• 3. What properties of carbon explain carbon’s ability to form many different macromolecules?

• 4. Critical Thinking Applying Concepts Explain why carbohydrates are polymers but lipids are not

Protein

• Amino Acids are the monomer links which contain C, H, N, and O. There are just over 20 different amino acids.

• Some proteins control the rate of reactions and regulate cell processes. Some are used to form bones and muscles. Others transport substances into or out of cells or help to fight disease.

• Amino acids are joined by peptide bonds (covalent bond between amino group (-NH2) and carboxyl group (-COOH)

• A protein can have four levels of structure: Primary, secondary (helix), tertiary, and quarternary structures

There are 20 Amino Acids

Dehydration Synthesis

• Homework SG 2-4

• all of SG chapter 2 should be complete

• Test chapter 2 coming up Friday???

1.  Key Concept Name four groups of organic compounds found in living things.

2. Key Concept Describe at least one function of each group of organic compounds.

3. What properties of carbon explain carbon’s ability to form many different macromolecules?

4. Critical Thinking Applying Concepts Explain why proteins are polymers but lipids are not.

Protein Structure

• Primary Structure– straight protein chain

• Secondary Structure – protein starts to bind to itself an kink the chain

• Tertiary Structure – Protein starts to wrap and fold around itself

• Quaternary Structure – Protein starts to fold and wrap around other proteins

Nucleic Acids

• Nucleotides are the monomer links which make up the polymer DNA or RNA

• Nucleic Acids store and transmit hereditary, or genetic, information

• Each nucleotide contains a

• 5-C sugar as either ribose or deoxyribose

• phosphate group (PO4-) and a

• Nitrogenous base (either: Adenine, Cytosine, Guanine, Thymine, or Uracil (only in RNA)

• DNA uses the sugar Deoxyribose

• RNA uses the sugar Ribose