chapter 2

Matter

• has mass and occupies space

  • 3 forms: solid, liquid, gas

Atom

• the smallest particle exhibiting chemical properties of an element

• composed of 3 subatomic particles (neutrons, protons, electrons)

Element

• 92 naturally occurring __ make up matter

• organized in periodic table

• usually occur as a mixture of isotopes

Most common elements of the human body (99% of weight)

most

1. Oxygen

2. Carbon

3. Hydrogen

4. Nitrogen

5. Calcium

6. Phosphorus

least

*Oh, Can Harry Not Cook Pasta

Minor elements of the human body (less than 1% of weight)

most

• Sulfur

• Potassium

• Sodium

• Chlorine

• Magnesium

• Iron

least

• SPS CMI

atomic mass unit (amu)

• a standard unit used to measure the mass of atoms and molecules

• 1 amu = one-twelfth (1/12) of the mass of a single carbon-12 atom.

Subatomic particles

• protons

• neutrons

• electrons

what differentiate these: mass & charge

Neutrons

• mass of 1 amu

• no charge

Protons

• mass of 1 amu

• positive charge of one (+1)

Electrons

• 0 amu

• 1/1800th of the mass of a proton or neutron

• negative charge of one (-1)

• located at varying distance from the nucleus in regions called orbitals (energy shells) or as an __ cloud

• # of __ is determined by atomic number

  • since all atoms are neutral, the # of negatively charged _ must sequal the # of positively charged protons

Chemical symbol

• unique to each element

• usually identified by first letter, or first letter + additional letter (C = carbon)

Atomic number

• number of protons in an atom of the element

• located above symbol name

• elements arranged by __ __ within rows

Average atomic mass

• mass of both protons + neutrons

• shown below the elements symbol on the table

Determining the # of subatomic particles

• proton # = atomic number

• neutron # = atomic mass - atomic number

  • neutron # = (p+n) - p

  • for ex, neutron # of Na = 23 - 11 = 12

• electron number = proton #

Cation

• positively (+) charged ions

  • formed by loosing electrons

    • from atoms with 1,2, or 3 electrons in their outer shell

• elements in columns IA, IIA, or IIIA, in periodic table (left side)

• “cats are pawsitive”

Anions

• negatively (-) charged ions

  • formed by gaining electrons

    • by atoms typically with 5, 6, or 7 electrons in the outer shell

• elements in columns VA, VIA, or VIIA (noble gases) in periodic table (right side)

Left side of periodic table

• loose /donate electrons (to form a completie valence shell)

• form cations when they react

Right side of periodic table

• gain electrons (to form a completie valence shell)

• form anions when they react

Anomalies in the period table

• iron (Fe)

  • it can form more than 1 type of ion, either a ferrous (Fe²+) or a ferric (Fe³+) ion

Shells

• an atom has __ of electrons surrounding the nucleus

• each has given energy level

• each holds a limited # of electrons

• innermost shell: 2 electrons

• second shell: up to 8 electrons

• shells close to the nucleus must be filled first

Isotopes

• different atoms of the same element

• same # of protons & electrons, diff # of neutrons

• identical chemical characteristics, diff atomic masses

  • for ex, carbon exists in 3 isotopes

  • carbon-12 with 6 neutrons

    • most prevalent type

  • carbon-13 with 7 neutrons

  • carbon-14 with 8 neutrons

• weighted average of atomic mass for all isotopes is the average atomic mass

Average atomic mass

• the weighted average of the atomic masses of its naturally occurring isotopes

Radioisotopes

• unstable forms of an isotope

• contain excess neutrons, so unstable

• radioactive: release radiation as they decay into a more stable isotope / loose nuclear components in the form of high energy radiation

  • alpha particles

  • beta particles

  • gamma rays

• may be traced & quantitfied by the high energy radiation these isotopes emit following their uptake into the body

scintigraphy

• nuclear medical imaging procedure whereby a body structures metabolic activity may be assessed with these isotopes

Physical half-life

• the time for 50% of radioisotope to become stable

  • the time it takes for a radioactive isotope to decay by half through nuclear processes,

• can vary from a few hours to thousands of years

Biological half-life

• the time required for half of the radioactive material from a test to be eliminated from the body

  • The time required for an organism, such as a human, to excrete or clear 50% of a specific substance.

Nonradioisotopes

• stable isotopes

• atoms of the same chemical element that have a different number of neutrons but do not break down or emit radiation over time

Radioisotopes introduced into the body during medical procedures

• Used by cells in a similar manner to nonradioisotopes

• Can trace products of metabolic reactions that use these elements

• Thyroid gland darker in areas where less radioactive iodine taken

up

• Can help locate a nodule

Periodic table organization

• organized horizontally across the table based on increasing atomic #

• organized into columns based on # of electrons in outer shell (valence shell)

  • elements tend to lose, gain, or share electrons to obtain complete outer shells with 8 electrons

• Column IA all have one electron in their valence shell

  • hydrogen, lithium, sodium, potassium

• each consecutive column has 1 additional electron in valence shell

• Column VIIA elements each have a full valence shell

  • results in chemical stability

  • helium, neon, etc, chemically inert noble gases

What two subatomic particles determine the mass of an atom?

• proton

• neutron

What two subatomic particles determine the charge of an atom

• proton

• electron

Uncharged atom

• has an equal number of protons and electrons

Octet rule

• atoms are most stable when they have 8 valence electrons (electrons in their outermost shell).

2-8-8 Rule

• electron-filling model for how electrons stack into shells (energy levels) around a nucleus

  • first shell holds up to 2 electrons, the second holds up to 8, and the third temporarily holds up to 8 before spilling over

Chemical compounds

• made of two or more different elements joined together by chemical bonds in a fixed ratio / stable associations between 2 or more elements combined in a fixed ratio

  • classified as ionic or molecular

Ionic compounds

• structures composed of ions held together in a lattice by ionic bonds/ electrostatic interactions

• ex

  • sodium chloride (NaCl)

  • sodium bicarbonate (NaHCO3)

  • magnesium chloride (MgCl2)

  • calcium phosphate Ca3(PO4)2

Ions

• atoms with a positive charge (cations, loose electrons) or a negative charge (anions, gain electrons)

• produced from loss or gain of 1 or more electrons

• significant physiological functions

• ex of common ones

  • sodium (Na^+)

  • calcium (Ca²+)

  • magnesium (Mg²+)

  • hydrogen (H^+)

  • chloride (Cl^-)

  • bicarbonate (HCO3^-)

  • phosphate (PO4 ³-)

  • potassium (K^+)

    • for ex, K^+ is used in sports drinks to replace the K^+ lost in sweat

      • K^+ in a large dose is used in some states for lethal injection

Sodium can reach stability by

• donating an electron

  • satisfies ocetet rule

  • now has 11 protons & 10 electrons

  • charge is +1 (cation)

Chlorine can reach stability by

• gaining an electron

  • satisfies octet rule

  • now has 17 protons & 18 electrons

  • charge is -1 (anion)

Polyatomic ions

• ions with more than 1 atom

  • for ex, bicarbonate ion (HCO3^-) & phosphate ion (PO4³-)

Intracellular

• the liquid found inside cells

• more negative

ex. potassium (K^+), phosphate (PO4³-), CO2, protein, magnesium (Mg²+),

“Poke Inside” “Kept inside” “Keep Most Positive” “KPMPC”

Extracellular

• the liquid found outside of cells

• more positive

• ex. sodium (Na^+), calcium (Ca²+), O2, chloride (Cl-), bicarbonate (HCO3^-)

• “NaCl is out in the ocean, and we are salty on the outside” / “NaCaClB” “NaClCa + bicarbonate

Common cations

• Sodium ion (Na^+)

• Potassium ion (K^+)

• Calcium Ion (Ca²+)

• Magnesium ion (Mg²+)

• Hydrogen ion (H^+)

“Silly Penquins Can Make Hugs”

Sodium ion (Na^+)

• most common extracellular cation

• particpant in conducting electrical signals in nerves & muscle

• most important in osmotic movement of water

• sodium gradient involved in cotransport of other substances across a plasma membrane

Potassium ion (K^+)

• most common intracellular cation

• participant in conducting electrical signals in nerves & muscle

• role in glycogen storage in liver & muscle

• function in pH balance

Calcium Ion (Ca²+)

• extracellular

• hardens bone & teeth

• muscle contraction

• exocytosis (including release of neurotransmitter)

• blood clotting

• second messenger in hormonal stimulation of cells

Magnesium ion (Mg²+)

• required for ATP production

Hydrogen ion (H^+)

• one proton + one electron

• concentration determines pH of blood & other fluids of the body

Common anions

• Chloride ion (Cl^-)

• Bicarbonate ion (HCO3^-)

• Phosphate ion (PO4)

Cheesy Burger Please

Chloride ion (Cl^-)

• alters nerve cell responsiveness to stimulation

• component of stomach acid (HCl)

• ___ shift in erythrocytes

Bicarbonate ion

• conversion of CO2 gas to HCO3^- which is transported in the blood

• buffering of pH in blood

Phosphate ion

• as Ca3 (PO4)2 it hardens bone & teeth

• component of phospholipids (membranes)

• component of nucleotides, including ATP & nucleic acids (DNA, RNA)

• most common intracellular anion

• intracelular buffer

Ionic bonds / Salt bridges

• complete transfer of electrons to create charged ions

• typically form between a metal & a nonmetal

• electrostatic attractions between + charged cation & - charged anions that form ionic compounds (lattice crystal/salt)

  • can form between an R group with a negative charge & an R group with a positive charge

  
  

• called salts

  • for ex. table salt (NaCL)

    • each sodium atom looses 1 outer shell electron to a chlorine atom

    • soidum & chlorine ions are held together by ionic bonds in a lattice crystal structure (ionic compound)

  • for ex. magnesium chloride (MgCL2)

    • each magnesium atom loses 1 electron to each of the 2 echlorine atoms

Whats the most common ionic compound in the body?

• calcium phosphate Ca3(PO4)

  • hardens bone & teeth

The elements that form the common ions of the human body (do not include the polyatomic ions

• Sodium (Na^+)

  • Role: Balances body fluids and helps send nerve impulses.

• Potassium (K^+)

  • Role: Regulates heart rhythm and supports muscle contraction.

• Calcium (Ca²+)

  • Role: Builds strong bones, teeth, and triggers muscle action.

• Magnesium (Mg²+)

  • Role: Supports immune function and helps produce energy.

• Chloride (Cl^-)

  • Role: Forms stomach acid and keeps body fluids balanced.

• Hydrogen (H^+)

  • Role: concentration dertermines pH of blood & other fluids of the body

Why ions form based on the octet rule

• Atoms form ions by gaining or losing valence electrons to reach a stable, filled outer shell with eight electrons

• Metals (Cations): Atoms on the left side of the periodic table have few valence electrons. They lose these electrons to become positively charged ions.

• Nonmetals (Anions): Atoms on the right side of the periodic table are close to having a full shell. They gain electrons to become negatively charged ions.

Could an ionic bond form between two cations or between two

anions?

• No, Ionic bonds require an electrostatic attraction between particles with opposite charges.

• Cations: These are positively charged ions.

• Anions: These are negatively charged ions.

• they push away from eachother instead of binding,

• * like charges repel

Covalently bonded molecule

• electrons shared between atoms of 2 or more diff elements

Molecular compound

• a chemical substance made of atoms that share electrons through covalent bonds

• molecules composed of 2 or more diff elements covalently bonded together

• ex. CO2, H2O

GICoS

• Give electron = Ionic Bond

• Covalent Bonds = Sharing electrons

Molecular formula

• indicates # & type of atoms within a molecule

• for ex. carbonic acid (H2CO3)

Structural formula

• indicates # & type of atoms within a molecule

• AND indicates arrangement of atoms within the molecule

• for ex. O=C=O (carbon dioxide)

• allows differentiation of isomers

  • same # & type of elements, but arranged differently in space

Isomer

• same # & type of elements/atoms, but arranged differently in space

• may have diff chemical properties

• ex. glucose. galactose, fructose

  • 6 carbon, 12 hydrogen, 6 oxygen

  • atoms arranged differently

Covalent bond

• form between nonmetals

• atoms share electrons (to fill outer shells) = resulting structure is molecule

• may be polar/nonpolar (shared unequally/equally)

  • determined by electronegativity

• occurs when both atoms require electrons to become chemically stable

• formed commonly in human body using HONC

  • are elements that most commonly form __ __ in the human body

• stabilize the folding of a protein

HONC 1,2,3,4

• Hydrogen (wants 1 bond)

• Oxygen (wants 2 bonds)

• Nitrogen (wants 3 bonds)

• Carbon (wants 4 bonds)

Single covalent bond

• one pair of electrons shared (2)

  • ex. between 2 hydrogen atoms

Double covalent bond

• two pairs of electrons shared (4)

  • ex. between 2 oxygen atoms

Triple covalent bond

• three pairs of electrons shared (6)

  • for ex, between 2 nitrogen atoms

Carbon Covalent Bonds

• wants 4 bonds

• Methane (CH4)

• Carbion dioxide (CO2)

• Ethanol (C2H5OH)

Carbon skeleton

• the arrangement of molecules that have numerous carbon atoms bonded together within their chemical structure

• the molecules “backbone”

• ex.

  • straight chain

  • branched chain

  • rings

Covalent Bonds Carbon skeleton formation

• bonds in straight chains, branched chains, or rings

• carbon present where lines meet at an angle

• additional atoms are only hydrogen

Electronegativity

• relative attraction of each atom for electrons

• determines how electrons are shared in covalent bonds

• determined both by the # of protons in the nucleus (pull on electrons) and the proximity of the valence electron shells to the nucleus

• in periodic table __ increases

  • from left to right across row

  • from bottom to top in column

  • * top right is most electronegative

• more ___ atom = develops a partial - charge

• less ___ atom = develops a partial + charge

  • written using Greek delta (δ) followed by superscript plus or minus

More electronegative atom

• develops a partial negative charge

Less electronegative atom

• develops a partial positive charge

Least to greatest electronegativity for 4 most common elements composing living organisms

hydrogen < carbon < nitrogen < oxygen
"Oh, No, Carl’s Here!"

"Have Copy Notes Only”

HNCO

Nonpolar molecules

• contain nonpolar covalent bonds / charges pull equally in all directions or if the bonds have no charge difference

  • however, ___ ___ may contain polar covalent bonds, if the polar covalent bonds cancel each other

    • ex. carbon dioxide

How to identify

• Same elements: Molecules made of only one pure element (like

O2 or N2) share electrons equally

Pure hydrocarbons: Molecules with only carbon and hydrogen (like methane,

CH4) are nonpolar because their atoms share pull evenly.

Symmetry: If the outer atoms are all identical and there are no extra unshared electrons on the middle atom, the pulls cancel each other out (like CO2)

Nonpolar covalent bond

• 2 atoms of same element have equal attraction for electrons

• If two different elements share a nearly identical pull on electrons, the charge stays balanced

• the bonded atoms have the same or very similar electronegativity values

  • difference is 0 to 0.4

• any bond with an electronegativity difference under 0.4 as nonpolar

• ex. C—C, C—H, O—-O bonds & oxygen (O2), triglyercide (fat), carbon dioxide (CO2)

Nonpolar examples

• C—C

• O—O

• C—H (the exception)

• oxygen (O2)

• triglyercide

• carbon dioxide (CO2) (

  • because its linear, symmetrical shape causes its individual bond pulls to cancel out

  • has polar covalent bonds that extend in opposite directions can be __ bc the partial charges cancel each other

• hydrocarbons

  • megthane gas (CH4)

Polar

• “Dipole”

• a molecule has a separation of electric charge, creating a positive side and a negative side like a tiny magnet

• refers to the different “poles” of partial electric charge

• “—-” opposites, north & south pole, WATER!!

Polar molecules

• contain polar covalent bonds / one side pulls harder on electrons than the other, creating a positive side and a negative side

How to identify

• Different outer atoms: If the atoms attached to the central atom are different, the pull is uneven (like CH3F)

• Lone pairs: If the central atom has leftover unshared pairs of electrons, it bends the shape and makes the molecule lopsided (like water, H2O or ammonia NH3)

• Asymmetry: If the shape is uneven, individual bond pulls do not cancel out

Polar covalent bond

• sharing of electrons unequally (between atoms of different elements)

• one atom pulls the shared electrons closer because it has a higher electronegativity

  • difference is 0.5 to 1.7

• creates a partial positive and partial negative charge on the molecule.

• ex. O—-H (H2O), C—O, N—H, N—O, glucose (C6H12O6), ethanol (C2H5OH), ammonia (NH3)

Polar examples

• O—H

• C—O

• N—H

• N—O

• water (H2O)

• glucose (C6H12O6)

• ethanol (C2H5OH)

• ammonia (NH3)

• carbohydrates

• O—H

• C—O

• N—H

• N—O

• water (H2O)

• glucose (C6H12O6)

• ethanol (C2H5OH)

• ammonia (NH3)

• carbohydrates

• hydroxl (-OH)

• carboxyl or carboxylic acid (—COOH)

• amine (-NH2)

• phosphate (PO4 ³-)

Amphipathic molecules

• large molecules w both polar & nonpolar regions

• form chemical barriers within the body, including membranes & micelles

• ex. phospholipids

Intermolecular attractions

• weak chemical attractions between molecules

• important for shape of complex molecules

• for ex, DNA & proteins

• unequal charges in nonpolar molecules

  • electrons orbiting nucleus briefly, unevenly distributed

  • induce unequal distribution of adjacent atoms of another nonpolar molecule

  • individually weak, collectively strong

Hydrogen bond

• forms between polar molecules (unequal sharing)

• electrostatic attraction between a partially positive hydrogen atom & a partially negative atom

• individually weak, collectively strong

• influences how water molecules behave

Hydrophobic interactions

• nonpolar molecules placed in a polar substance

• if occurring between parts of large molecule, termed intra molecular attraction

Assign the partial charges between nitrogen and hydrogen (N—H) in a polar covalent bond.

Nitrogen (N): (partial negative)

• Nitrogen is highly electronegative (3.04).

Hydrogen (H): (partial positive)

• Hydrogen has lower electronegativity (2.20).Nitrogen pulls shared electrons closer.

Identify the exception to the rule that polar covalent bonds are formed between two different types of atoms.

• The carbon-hydrogen (C–H) bond

• even though they are two different elements, their electronegativity difference is extremely small (about 0.35)

• bc this difference is so tiny, they pull on electrons with nearly equal strength

• as a result, the C–H bond shares electrons fairly evenly and is treated as a nonpolar covalent bond

Are O2 and CO2 nonpolar or polar molecules? Explain why.

O2: nonpolar

• It is a diatomic molecule made of two identical oxygen atoms. Because both atoms have the exact same electronegativity (ability to attract electrons), they share the bonding electrons equally.

CO2: nonpolar

• Oxygen is more electronegative (better at attracting electrons) than carbon. Each double bond pulls shared electrons closer to the oxygen atom, creating an individual "bond dipole" (a pull of electrical charge)

Organic molecules

• carbon + hydrogen

• are (or were) part of a living organism

  • ex. glucose, proteins, triglyceridedes

Inorganic molecules

• all other molecules

  • ex. water, salts, acids, bases

Water

• is polar

  • 1 oxygen atom bonded to 2 hydrogen atoms

  • oxygen atom has 2 partial negative charges

  • each hydrogen has 1 single partial positive charge

  • can form 4 hydrogen bonds with other water molecules

• inorganic

• universal solvent (esp of the body)

• bathes organs in fluid which provides protection & support

• plays a major role in bodies chemical reactions

• neutral

  • + hydrogen ions (H^+) = - hydroxide ions (-OH)

• 2/3 of body weight

Phases of water

• depending on temp

• gas (water vapor)

  • substances with low molecular mass

• liquid (water)

  • almost all water in the body

  • liquid at room temp due to hydrogen bonding

• sold (ice)

Functions of liquid water

• transports

  • substances dissolved in water move easily throughout body

• lubricates

  • decreases friction between body structures

• cushions

  • absorbs sudden force of body movements

• excretes wastes

  • unwanted substances dissolve in water are easily eliminated

Properties of water

• cohesion

  • attraction between water molecules due to hydrogen bonding

• surface tension

  • inward pulling of cohesive forces at surface of water / pull of water molecules at the surface

  • causes moist sacs of air in lungs to collapse

    • surfactant, a lipoprotein, prevents collapse

  • adhesion

    • attraction between water molecule & a substance other than water

  • high specific heat

    • • Amount of energy required to increase temperature of 1 gram of a

      substance by 1 degree Celsius

      • Water’s value extremely high due to energy needed to break

      hydrogen bonds

      • Contributes to keeping body temperature constant

  • high heat of vaporization

  • • Heat required for release of molecules from a liquid phase into a

    gaseous phase for 1 gram of a substance

    • Water’s value very high due to hydrogen bonding

    • Sweating cools body

    • Excess heat dissipated as water evapora

Surface tension

• inward pulling of cohesive forces at surface of water / pull of water molecules at the surface

• tendency of a liquids surface to resist breaking when placed under stress

  • bc at the surface water molecules are pulled by hydrogen bonds in only 3 directions whereas water molecules that are internal in the liquid are pulled by hydrogen bonds in 4 directions

• causes moist sacs of air in lungs to collapse

  • surfactant, a lipoprotein, prevents collapse

  • contributes to the need to produce surfactant & prevent collapse of the alveoli

surfactant

• a substance that lowers the surface tension of a liquid, allowing it to spread easily and mix with other substances like oil and water

Cohesion

• attraction between water molecules due to hydrogen bonding

• attraction of water molecules to other water molecules

Adhesion

• attraction between water molecule & a substance other than water

Temperature

• measure of kinetic energy of atoms or molecules within a substance

• higher __ = greater kinetic energy