A&P CH2 chem

Matter

1. Matter is anything that has mass and takes up space
a. Mass is the amount of matter a substance contains & remains constant wherever the object is
b. Weight is the force of gravity acting on a mass
2. Matter exists in 3 forms:
a. Solid: definite shape and definite volume
b. Liquid: changeable shape; definite volume
c. Gas: changeable shape and changeable volume
3.All forms of matter are composed of chemical elements

Chemistry

- Is about matter and energy
- the science of structure and interactions of matter
- the interaction of matter requires or generates energy

Energy

- Energy is the capacity to do work or put matter into motion
- Energy does not have mass, nor does it take up space
- he greater the work done, the more energy is used doing i

Chemistry and Physiological Reactions

Body is made up of many chemicals
•Chemistry underlies ALL physiological processes:
•All cellular functions, movement, digestion, pumping of heart, nervous system
•Chemistry can be broken down into: basic Chemistry and biochemistry

Basic chemistry

Organic (carbon-based molecules)
•inorganic

Biochemistry

•organic and inorganic chemistry
•All chemical reactions that take place in biological systems

Forms of energy

•Chemical energy
•Electrical energy
•Mechanical energy: Directly involved in moving matter
•Radiant or electromagnetic energy

chemical elements

- All matter is composed of elements
- Elements are substances that cannot be broken down into simpler substances by ordinary chemical methods
Elements are given letters as chemical symbols
- Four elements make up 96% of body:
•O \= oxygen
•C \= carbon
•H \= hydrogen
•N \= nitrogen
- 9 elements make up 3.9% of body
- 11 elements make up

Atoms

•Atoms are the units of matter
•Chemical elements are composed of (units of matter) atoms of the same type
•Atoms are the smallest units of matter that retain the properties and characteristics of an element
•Atoms have subatomic particles
•Two models: electron cloud model and electron shell model

Atomic number and mass number

Atoms of different elements differ in the number of subatomic particles
•Atomic number \= number of protons in the nucleus of an atom. written as a subscript to the left of its atomic symbol.
•Mass number \= number of protons + number of neutrons in an atom
•Number of electrons \= number of protons → neutral (no charge)

Atomic mass

Atomic mass assumes the mass of a:
•Neutron \= 1.008 daltons
•Proton \= 1.007 daltons
•Electron \= 0.0005 daltons
Atomic mass \= neutrons + protons
the size of all molecules in physiology, biochemistry and O Chem is given in Daltons or kDa
The atomic mass/weight of an element is the average mass of all its naturally occurring isotopes

Isotope

Structural variations of same element
•Atoms contain same number of protons but differ in the number of neutrons they contain; atomic numbers are same, but mass numbers different... They are HEAVIER!
- iso \= same; topos \= place

Atomic weight

Average of the mass numbers of all isotopes

Radioisotopes

- An isotope that has an unstable nucleus and undergoes radioactive decay
- heavier isotopes of many elements are unstable, and their atoms decompose spontaneously into more stable forms
- disintegration of a radioactive nucleus may be compared to a tiny explosion. It occurs when subatomic alpha (α) particles (packets of 2p+2n2p+2n), beta (β) particles (electron-like particles), or gamma (γ) rays (electromagnetic energy) are ejected from the atomic nucleus
- dense nuclear particles are composed of even smaller particles called quarks that associate in one way to form protons and in another way to form neutrons. The "glue" that holds these nuclear particles together is weaker in the heavier isotopes. When radioisotopes disintegrate, the element usually transforms to a different element.
- Radioisotopes are a valuable tool for biological research and medicine
•Share same chemistry as their stable isotopes so will be taken up by body
•Can then be used for diagnosis of disease
•All radioactivity can damage living tissue
•Some types can be used to destroy localized cancers
•Some types cause cancer (Radon from uranium decay causes lung cancer)

Electrons

Electrons occupy areas around nucleus called electron shells
•Each shell contains electrons that have a certain amount of kinetic and potential energy, so shells are also referred to as energy levels (electron shell & energy level used interchangeably)
•Depending on its size, an atom can have up to 7 electron shells
•Shells can hold only a specific number of electrons; the shell closest to nucleus is filled first
•Shell 1 can hold only 2 electrons
•Shell 2 holds a maximum of 8 electrons
•Shell 3 holds a maximum of 18 electrons
•Outermost electron shell is called valence shell
•Electrons in valence shell have the most potential energy because they are farthest from nucleus (energy they absorbed to overcome the nuclear attraction and reach the more distant energy levels) more likely to interact with other atoms bc they are least tightly held by their own nucleus
- electron equal in charge to proton but only has 1/2000 mass of proton (0 amu)

Ion

An atom or group of atoms that has a positive or negative charge. (Gained or lost an electron)

Molecule

two or more atoms held together by covalent bonds (sharing electrons)

Compound

a molecule made up of 2 or more different elements
•It can be broken down into 2 or more different elements
•Compounds are formed by chemical bonds

Chemical bonds

A chemical bond occurs when atoms are held together by forces of attraction
•The number of electrons in the valence shell determines the likelihood that an atom will form a chemical bond with another atom
•The fate of these electrons determines the type of bond that is formed
•Types of bonds:
•Covalent
•Ionic
•Hydrogen bonds
•Van Der waals forces

Ionic bonds

Number of protons does not equal number of electrons
•Ionic bonds involve the transfer of valence shell electrons from one atom to another, resulting in ions
•One becomes an anion (negative charge)\= Atom that gained one or more electrons\= electron acceptor
•One becomes a cation (positive charge)\= Atom that lost one or more electrons\= electron donor

Covalent bonds

formed by sharing of two or more valence shell electrons between two atoms
•Sharing of 2 electrons results in a single bond
•Sharing of 4 electrons is a double bond
•Sharing of 6 electrons is a triple bond
•Allows each atom to fill its valence shell at least part of the time
•Two types of covalent bonds:
•Polar and nonpolar covalent bonds
•Nonpolar covalent bonds
•Equal sharing of electrons between atoms
•Results in electrically balanced, nonpolar molecules such as CO2

covalent bond

Sharing of electrons within a molecule
•Single, double, and triple bonds
- strongest bond

polar covalent bond

Unequal sharing of electrons between 2 atoms
•Results in electrically polar molecules
•Atoms have different electron-attracting abilities, leading to unequal sharing
•Atoms with greater electron-attracting ability are electronegative (have 6 or 7 ve) , and those with less (one or two valence electrons) are electropositive: electron-attracting ability is so low that they usually lose their valence shell electrons to other atoms
•H2O is a polar molecule
•Oxygen is more electronegative, so it exerts a greater pull on shared electrons, giving it a partial negative charge and giving H a partial positive charge
•Having two different charges is referred to as dipole
•This Dipoles involving H permit the formation of another type of bonds: Hydrogen bonds

hydrogen bond

result from attraction of oppositely, partially charged parts of molecules... These molecules have a dipole moment
When in bulk, these molecules form a lattice where molecules are connected to each of its neighbors via H-bonds
Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule
Not true bond, more of a weak magnetic attraction
Common between dipoles such as water
What makes water liquid
Also act as intramolecular bonds, holding a large molecule in a three-dimensional shape

Cohesion

Hydrogen bonds between water molecules give water cohesion
•Cohesion is the tendency of like particles to stay together

surface tension

Hydrogen bonds create surface tension
•Surface tension is a measure of the difficulty of stretching or breaking the surface of a liquid
•Cohesion and surface tension give water its bulk properties, which contribute to:
•Phase separation or compartmentalization
•Self-organization of macromolecules
•Hence we talk about hydrophilic vs hydrophobic molecules to define the type of interaction between molecules and water

Mixtures

Most matter exists as mixtures: two or more components that are physically intermixed
•Three basic types of mixtures
•Solutions
•Colloids
•Suspensions

Solutions

Are homogeneous mixtures, meaning particles are evenly distributed throughout
True solutions are usually transparent
•Example: air (gas solution), salt solution, sugar solution
•Most solutions in body are true solutions of gases, liquids, or solids dissolved in water

Colloids

Also known as emulsions; are heterogeneous mixtures, meaning that particles are not evenly distributed throughout mixture
do not settle out. However, they do scatter light

Concentration of true solutions

Three common ways to express concentrations:
1.Percent of solute in total solution
•How many parts of solute are in 100 total parts of solution
•Solvent is usually water
•Example: 10 parts salt to 90 parts water is a 10% salt solution
2.Milligrams per deciliter (mg/dl)
•Deciliter equals 1/10th of a liter, i.e 1dl \= 100ml
•Example: normal fasting blood glucose levels are around 80 mg/dl
3.Molarity (M) is number of moles of solute per liter of solvent (water)
•1 mole of a compound is equal to its molecular weight (sum of atomic weights) in grams
•Example: glucose (C6H12O6) has a molecular wt of 180.12 amu, so 180.12 grams of glucose added to enough H2O to make 1 liter is a 1 M solution of glucose
•1 mole of any substance always contains 6.02 × 1023 molecules

Solvent

substance present in greatest amount
•Usually a liquid, such as water

percentage and molarity

Solute

substance dissolved in solvent
•Present in smaller amounts
•Example: blood sugar - glucose is solute, and blood (plasma) is solvent

colloids

- Also known as emulsions; are heterogeneous mixtures, meaning that particles are not evenly distributed throughout mixture
• Can see large solute particles in solution, but these do not settle out
• Gives solution a cloudy or milky look
• Some undergo sol-gel (solution to gel) transformations
• Example: Jell-O goes from liquid to gel
• Cytosol of cell is also a sol-gel type solution
- sol-gel transformations underlie many important cell activities, such as cell division and changes in cell shape

suspensions

- Heterogeneous mixtures that contain large, visible solutes that do settle out
• Example: mixture of water and sand
• Blood is considered a suspension because if left in a tube, the blood cells will settle out

concentration of true solutions

- three common ways to express concentrations:
1. Percent of solute in total solution
2. Milligrams per deciliter (mg/dl)
3. Molarity (M) is number of moles of solute per liter of solvent (water)

1. Percent of solute in total solution

• How many parts of solute are in 100 total parts of solution
• Solvent is usually water
• Example: 10 parts salt to 90 parts water is a 10% salt solution

2. Milligrams per deciliter (mg/dl)

• Deciliter equals 1/10th of a liter, i.e 1dl \= 100ml
• Example: normal fasting blood glucose levels are around 80 mg/dl

3. Molarity (M) is number of moles of solute per liter of solvent (water)

• 1 mole of a compound is equal to its molecular weight (sum of atomic weights) in grams
• Example: glucose (C6H12O6) has a molecular wt of 180.12 amu, so 180.12 grams of glucose added to enough H2O to make 1 liter is a 1 M solution of glucose
• 1 mole of any substance always contains 6.02 × 1023 molecules of that substance
• This number is called Avogadro's number
• Molarities in the body are so small (can be 0.0001 M), they are expressed in millimoles (mM) so1000 mM 1 M

compounds

• Compounds are represented as molecular formulas
• Example: H2O or C6H12O6 or H2 or CH4
• In chemical equations, subscripts indicate how many atoms are joined by bonds, whereas prefix means number of unjoined atoms (example: 4H)

chemical reactions

- occur when new bonds are formed or old bonds are broken
Reactants - starting substances
Products - ending substances

metabolism

The totality of all chemical reactions in the body

types of chemical reactions in body

1. Synthesis
2. Decomposition
3. Exchange
4. Reversible
5. Oxidation-reduction

oxidation reduction reaction

- These reactions transfer electrons between atoms and molecules and always occur in parallel (when one substance is oxidized another is reduced)
- Oxidation - loss of electrons and energy release
- Reduction - gain of electrons and energy gain
- redox reactions happen when ionic bonds form

synthesis reaction

- (combination) reactions involve atoms or molecules combining to form larger, more complex molecule
• Used in anabolic (building) processes

decomposition

- reactions involve breakdown of a molecule into smaller molecules or its constituent atoms (reverse of synthesis reactions)
• Involve catabolic (bond-breaking) reactions

exchange reaction

AB + CD --\> AD + CB
- also called displacement reactions, involve both synthesis and decomposition• Bonds are both made and broken

energy and chemical reactions

Chemical reactions either release or consume energy
Energy is the capacity to do work; it can be in various forms:
Potential
Kinetic
Chemical
Thermal
Mechanical
Law of conservation of energy - energy can neither be created nor destroyed but it can be converted from one form to another
Your body is constantly converting one form of energy to another to maintain Homeostasis... when this stops you reach a low energy state of equilibrium: death

energy transfer

Exergonic reactions (Release E): products have less energy than reactants; catabolic and oxidative reactions are exergonic.
Endergonic reactions (Consume E) Endergonic reactions are used to STORE Energy, for later use; contain more potential energy in their chemical bonds than did the reactants; anabolic
Activation energy: required for a chemical reaction to take place

Catalysts

- allow reactions to proceed by lowering the activation energy
- Catalysts do not change the total energy of the reaction
- Catalysts speed the rate of the reactions
- They do NOT add energy to endergonic reactions (or store E from exergonic)
- Therefore, the body's metabolism requires the coupling of chemical reactions
- An endergonic reaction is coupled to an exergonic reaction

inorganic compounds

- usually lack carbon and are simple molecules
- Ions, ionic salts, elemental molecules
- Water is the most important and abundant inorganic compound in all living things

organic compounds

always contain C & H, usually contain O, and always have covalent bonds

Water as a polar molecule

- Water is the ideal medium
- Water has a:High heat capacity &High heat of vaporization
- Water is a major component of our body fluids and helps reduce friction as membranes and organs slide over one another (cushioning) cerebrospinal fluid surrounding the brain exemplifies water's cushioning role.
- it has high surface tension

Water in chemical reactions

- Water CAN participate in chemical reactions
- In a hydrolysis reaction water is added to break bonds
- In a dehydration reaction water is removed to make bonds. This is also called a synthesis reaction

acids, bases, and salts

- Mixtures containing ions as particles (solutes) in water (solvent) can form an acid, a base or a salt.
- The ions dissociate and are surrounded by water molecules
- If if sheds H+ or OH- ions it can change the pH of the solution; it changes the concentration of H+ found in the solution →pH

pH

- The pH is a scale to measure the concentration of H+ in solution; it is the inverse log of the [H+].
- more hydrogen ions in a solution, the more acidic the solution is. Conversely, the greater the concentration of hydroxyl ions (the lower the concentration of H+H+), the more basic, or alkaline solution is
- devised by a Danish biochemist and beer brewer named Sören Sörensen in 1909
- pH scale 0-14 & is logarithmic: each successive change of one pH unit represents a tenfold change in hydrogen ion concentration
- most acidic at 10^0 [H+] \= 10^-14 [OH-]
- most basic at 10^-14 [H+] \= 10^0 [OH-]
- pH of 7 (at which [H+]H+ is 10−710−7M), the solution is neutral—neither acidic nor basic; same amount oh hydrogen and hydroxyl ions

buffer systems

- Maintaining blood and Interstitial fluid pH near 7.4 is CRITICAL for the human body
- Buffers are chemicals that regulate pH by donating or removing H+ from the solution
- Buffers can do this because they consist of a combination of a weak acid and a corresponding weak base
- Maintenance of pH is fundamental for body fluid homeostasis, i.e. cellular function and survival
- Buffer systems help to regulate pH by converting strong acids or bases into weak acids or bases
- pH below 7 are acidic, above 7\= alkaline

chemical energy

Stored in bonds of chemical substances, when chemical reactions occur, atoms become rearranged and P energy turns into K energy

electrical energy

Results from movement of charged particles.
electrical currents are generated when charged particles (ions) move along or across cell membranes. The nervous system uses electrical currents, called nerve impulses (or action potentials), to transmit messages from one part of the body to another

radiant/electromagnetic energy

- energy that travels in waves. These waves, which vary in length, are collectively called the electromagnetic spectrum
- Light energy, which stimulates the retinas of our eyes, is important in vision. Ultraviolet waves cause sunburn, but they also stimulate your body to make vitamin D

converting forms of energy

- chemical energy (in gasoline) that powers the motor of a speedboat is converted into the mechanical energy of the whirling propeller that makes the boat skim across the water
- energy conversions inefficient bc some energy is always released as heat. ex: electrical energy turns into light energy with lightbulb but lightbulb is hot
- all energy conversions in the body liberate heat. This heat helps to maintain our relatively high body temperature, which influences body functioning
- higher temp, faster body's chemical reactions occur

physical properties

- those we can detect with our senses (such as color and texture) or measure (such as boiling point and freezing point)

chemical properties

- the way atoms interact with other atoms (bonding behavior) and account for the facts that iron rusts, animals can digest their food,

Planetary Model (Bohr)

- electrons move around the nucleus in fixed, circular orbits
- easiest to depict so used in illustrations althouh orbital model more accurate

orbitals

- regions around the nucleus in which a given electron or electron pair is likely to be found most of the time. This more modern orbital model is more useful for predicting the chemical behavior of atoms
- depicts probable regions of greatest electron density by denser shading (this haze is called the electron cloud)

half life

time required for a radioisotope to lose one-half of its activity is called its half-life. The half-lives of radioisotopes vary dramatically from hours to thousands of years

Compounds vs. Mixtures

- no chemical bonding occurs between the components of a mixture. The properties of atoms and molecules are not changed when they become part of a mixture
- mixture can be separated by physical means (straining, evaporation) compounds can only be separated by chemical means (breaking bonds)

chemically inert

elements with a full valence shell or 8 outer electrons
means they are unreactive
ex: noble gases

octet rule

atoms react by gaining or losing electrons so as to acquire the stable electron structure of a noble gas, \= eight valence electrons

intramolecular bond

(bonds within molecules), which hold different parts of a single large molecule in a specific three-dimensional shape. Some large biological molecules, such as proteins and DNA, have numerous hydrogen bonds that help maintain and stabilize their structures.

relative proportion

Balanced equations indicate the relative proportion of each reactant and product.

chemical equilibrium

In a chemical reaction, the state in which the rate of the forward reaction equals the rate of the reverse reaction, so that the relative concentrations of the reactants and products do not change with time.

factors that affect chemical reactions

- Temperature. Higher temperatures increase the kinetic energy of particles and the force of their collisions, increasing the rate of chemical reactions
- Concentration: High concentrations of reacting particles increase the chances of successful collisions, and reactions progress faster. unless products or reactants are added/removed, chemical equilibrium will eventually occur
- Particle size. The smaller the reacting particles, the faster the chemical reaction. Smaller particles move faster than larger ones (collide more frequently and forcefully)
- Catalysts: substances that increase the rate of chemical reactions without themselves becoming chemically changed or part of the product. Biological catalysts are called enzymes

molecular chaperones

enzymes that aid in the desired folding of proteins.

Prostaglandins

purines

Bases with a double-ring structure.
Adenine and Guanine

An organic compound is analyzed, and it has twice as many hydrogen atoms as oxygen atoms. This compound is most likely a

Carbohydrates have CHO with a 1:2:1 ratio

enzymes

increase the rate of chemical reactions by lowering activation energy

biochemistry

- is the study of chemical composition and reactions of living matter•
All chemicals either organic or inorganic Both equally essential for life

inorganic compounds

- Water, salts, and many acids and bases
• Do not contain carbon

organic compounds

- Carbohydrates, fats, proteins, and nucleic acids
• Contain carbon, are usually large, and are covalently bonded

organic molecules

- organic compounds always contain carbon
- EXCEPT for CO2 and CO which are inorganic

carbon

- is electroneutral: shares electrons; never gains or loses them, forms four covalent bonds with other elements, carbon is unique to living systems
- carbons can combine in a variety of shapes
- carbon atoms combine with other elements (H, O,S, P, etc) to from functional groups
- these FUNCTIONAL GROUPS become the site of reactions
- carbon compounds dont dissolve easily in water
- carbon compounds are a good source of energy

functional groups in organic compounds

- Functional groups in organic molecules are reactive sites of the molecule:
Determine their interaction with water
Participate in chemical reactions
C-C bonds are non-reactive & hard to break
organic molecules are very large molecules, but their interactions with other molecules typically involve only small, reactive parts of their structure \= functional groups

hydroxyl (major functional group)

R-O-H
- alcohols contain an -OH group which is polar & hydrophilic due the electronegative oxygen atom
- molecules with many -OH groups dissolve easy in water

sulfhydryl (major functional group)

R-S-H
thiols have an -SH group
- polar and hydrophilic due to its electronegative S atom
- certain amino acids (like cysteine) contain -SH groups which help stabilize shape of proteins

carbonyl (major functional group)

carbon double bonded to an oxygen
R-C-R & R-C-H
ketones contain a carbonyl group within the carbon skeleton
polar & hydrophilic due to its electronegative O atom
aldehydes have a carbonyl group at the end of the carbon skeleton

carboxyl (major functional group)

contain a carboxyl group at the end of the carbon skeleton
all amino acids have a -COOH group at one end
the negatively charged form predominates at the pH of body cells & is hydrophillic

major organic compounds

-carbohydrates
- lipids,
- proteins
- nucleic acids
many are polymers

polymers

• Chains of similar units called monomers (building blocks)
• Synthesized by dehydration synthesis
• Broken down by hydrolysis reactions

dehydration synthesis and hydrolysis

• Polymeric organic compounds (and other macromolecules) are synthetized via dehydration reactions, which remove a H molecule from one monomer and OH from second monomer
• polymeric organic compounds are broken down by hydrolysis reactions, which add a H2O molecule between two polymeric units (add back H to one and OH to another

carbohydrates

- sugars and starch
- Contain C, H, and O; Hydrogen and oxygen are in 2:1 ratio
- represent 1-2% of cell mass.
- provide most of energy needed for life
can be:
1. Simple saccharide (monomer)
2. Di- saccharide (dimer)
3. Poly- saccharide (polymer)

Monosaccharides

glucose, fructose, galactose, deoxyribose (in DNA), ribose (RNA)
pentose (five-carbon) and hexose (six-carbon) sugars
single-chain or single-ring structures containing from three to seven carbon atoms
- fructose & galactose isomers of glucose they have the same molecular formula (C6H12O6), but their atoms are arranged differently, giving them different chemical properties
water molecule is added as each bond is broken

Dissacharides

sucrose \= glucose + fructose
lactose\= glucose + galactose
maltose \= glucose + glucose
Disaccharides are too large to be transported through cell membranes, so they must be hydrolyzed to monosaccharides before being digested

Polysaccharides

- glycogen: stored form of carbs in animals
- startch: stored form of carbs in plants & main carb in food
- cellulose: part of cell wall cant be digested by humans but can aid in digestion bc provides bulk (fiber) to move feces
- polymers of simple sugars linked together by dehydration synthesis. Because polysaccharides are large, fairly insoluble molecules, they are ideal storage products.

Monosaccharides

• Simple sugars containing three to seven carbon atoms
• (CH2O)n: general formula, n \= number of carbon atoms
• Monomers of carbohydrates
• Important monosaccharides
• Pentose sugars
• Ribose and deoxyribose
• Hexose sugars
• Glucose (blood sugar)

Disaccharides

• Double sugars
• Too large to pass through cell membranes•
Important disaccharides
• Sucrose, maltose, lactose
• Formed by dehydration synthesis of two monosaccharides
• glucose + fructose → sucrose + water

Polysaccharides

• Polymers of monosaccharides
• Formed by dehydration synthesis of many monomers
• Important polysaccharides
• Starch: carbohydrate storage form used by plants
• Glycogen: carbohydrate storage form used by animals
• Not very soluble

lipids

• Contain C, H, O, but less than in carbohydrates, and sometimes contain P (phospholipids)
• Insoluble in water
• Main types: Triglycerides, Phospholipids, Steroids, Eicosanoids

saturated fatty acids

All carbons are linked via single covalent bonds, resulting in a molecule with the maximum number of H atoms it can hold (saturated with H)
These create LINEAR molecules which can pack closely together forming a solid at room temperature (Example: animal fats, butter)