Chapter 2: The Nature of Molecules and the Properties of Water Flashcards
Matter
any substance that has mass and occupies space
Mass
the amount of a substance
Atom
small particles that make up matter
Weight
the force gravity exerts on a substance
Atomic number
number of protons in an atom
Atomic mass/weight
sum of the masses of its protons and neutrons
What are the masses of atoms and subatomic particles measured in?
Daltons
Cation
atom with more protons than electrons with net + charge
Anion
atom with less protons than electrons with net - charge
Ion
- atoms where the number of protons is not the same as the number of electrons
- gain or loss of electrons
Proton
- positive
- found in nucleus
- 1 dalton
Neutron
- neutral
- found in nucleus
- 1 dalton
Electron
- negative
- orbit nucleus in energy levels
- negligible mass
Isotope
atoms of a single element with different number of neutrons
ex) Carbon-12 (stable) , Carbon-13 (stable) , Carbon-14 (radioactive) - used to determine ages of fossils
Radioactive decay
nucleus of an atom breaks up into elements with lower atomic numbers and releases energy
Orbital
regions where electrons are likely located
Energy levels
- outer levels have more energy
- 1st level with 1 orbital: holds 2 electrons
- 2nd level with 4 orbitals: holds 8 electrons
- 3rd level can hold 18 electrons
Chemical behavior
- based on valence electrons
- complete valence shell = nonreactive
- incomplete valence shell = chemically reactive
Octet rule
atoms tend to establish completely full outer energy levels
Molecule
- made up of two or more atoms of the same or different elements bonded together
- ex) O2
Compound
- two or more different elements bonded together
- ex) H2O
What four elements make up 96% of living matter?
carbon, oxygen, hydrogen, and nitrogen
Periodic table trends
- elements up and down have same number of electrons in outer shells
- atoms are electrically neutral
Chemical bonds
join molecules together
Covalent bond
- two atoms SHARE one or more pair of valence electrons
- can be single (share a pair of electrons) , double (share two pairs of electrons) , or triple (share three pairs of electrons)
- form stable molecules
- strength increases with number of shared electrons
Ionic bond
- electrons are TRANSFERRED from one atom to another
- attraction of opposite charges
- form crystals
Hydrogen bond
- weak attraction between molecules or parts of the same molecule
- slightly positive hydrogen in one molecule attracted to slightly negative atom in another molecule
- responsible for emergent properties of water
- ex) electrons spend more time orbiting oxygen than hydrogens, so oxygen becomes slightly negative
Van der Waals Interactions
- attractions between every changing + and - “hot spots” in covalently bonded nonpolar molecules
Polar covalent bonds
- unequal sharing of electrons
- regions of partial negative charge near more electronegative region
Nonpolar covalent bonds
- equal sharing of electrons
- ex) methane
Electronegativity
- atoms differ in affinity for electrons
- usually increases left to right across periodic table
- usually decreases down the column of periodic table
Oxidation
the loss of electrons (or loss of hydrogen atoms)
Reduction
the gain of electrons (or gain of hydrogen atoms)
Chemical reactions
- make and break chemical bonds
- many are reversible
- matter is conserved
- reactions must be balanced
Equilibrium
- rate of formation of products = rate of breakdown of products
- no change in concentrations
What does the shape of a molecule tell us?
- molecules with similar shapes have similar functions
- 2 atoms = linear shape
- water is shaped like a V
Catalyst
substance that increases reaction rate
Cohesion
- polarity of water which makes water molecules attracted to each other
- ex) water droplets bead up
Adhesive
- polarity of water which makes it attracted to other polar molecules
- ex) meniscus in graduated cylinder; wet microscope slides stick together
Capillary action
- liquid moves up through narrow tube
- results from cohesion and adhesion
- ex) water goes from roots to leaves; drinking straws; paper towel wicking
Surface tension
- caused by cohesion
- allows insects to walk on water
Density of water
- ice is less dense than liquid water because hydrogen bonds keep the molecules further apart
- ex) lakes freezing from top down
Kinetic energy
energy of motion
Heat
total amount of energy in a system
Temperature
average amount of kinetic energy in molecules
Calorie
amount of heat required to raise temp of water by 1 degree Celsius
Specific heat
- amount of heat a substance needs for a given increase of temperature
- water has a high specific heat
- hydrogen bonds absorb heat when they break, release it when they form
- takes a lot of energy to raise 1 g of water 1 degree C because hydrogen bonds need to break
- heats up slowly and retains heat for longer
- can absorb lots of heat, but with minimal temperature changes
- allows for homeostasis
How are specific heat and polarity related?
more polar = higher specific heat
High Heat of Vaporization
- takes a lot of energy to convert liquid water to vapor
- many bonds must be broken for water to evaporate
- ex) used by organisms when sweating to cool off
Hydrophobic
- nonpolar covalent molecules
- water fearing, don’t dissolve in water
- ex) fats, oils
Hydrophilic
- polar covalent molecules
- water loving
- form hydrogem bonds with water
Amphipathic
- molecules with polar and nonpolar regions
- ex) phospholipids: polar head/nonpolar tail
Mole
weight of a substance in grams that relates to the atomic masses of all the atoms in a molecule of the substance
Ionization
covalent bonds within water may break and form ions
pH
- measures the degree of acidity
- measures hydrogen ion concentration of a solution
- water pH = 7 = neutral
Difference on pH scale
pH 2 has 1000 times more H+ ions than pH 5
Acid
- substance that causes an increase in H+ ions
- pH is less than 7
Base (alkaline)
- substance that causes increase in OH- ions
- pH is more than 7
Buffer
- substance that resists changes in pH
- can donate H+ to solutions or remove them
- maintain homeostasis
Bicarbonate Buffer System
- HCO3- = bicarbonate (weak base)
- H2CO3 = carbonic acid (weak acid)
- helps to balance blood pH
Does CO2 increase or decrease pH?
decrease, becomes more acidic
Acidosis and its solution
- when the blood pH becomes more acidic and can harm kidneys
- solution: patients are given HCO3- to increase pH and decrease H+
How many of the 90 naturally occurring elements on Earth are found in living systems?
12
Why are the noble gases more stable than other elements of the periodic table?
all eight electrons fill up their outer energy level, so they are inert (nonreactive)
List the bonds in order form weakest to strongest
van der Waals, hydrogen bonds, ionic bonds, covalent bonds
What factors influence the extent of chemical reactions
1.) Temperature - it increases the reaction rate because the reactants bump into each other at greater frequency
2.) Concentration of reactants and products - more reactants speed up reactions, BUT more products slows down reaction unless it is reversible
3.) Catalysts - speed up reactions; example is enzyme
How much of the body is composed of water?
2/3
What is water’s polarity and bonding like?
- neutral in charge and polar covalent
- covalent bonds between hydrogen and oxygen atoms result in uneven sharing of electrons
- polar because hydrogen atoms are slightly positive, while oxygen area is slightly negative
Polar molecules
- molecule has uneven pattern of electric charge
- results in hydrogen bonding between water molecules
- responsible for emergent properties of water
Why is water a “universal” solvent?
- dissolves many polar substances
- separates ionic substances into ions
- provides medium for molecules to interact
What substances dissolve in water?
- ionic substances, which dissolve into ions
- polar covalent substances
Solution
uniform mixture of molecules of 2 or more substances
Solvent
substance present in the greatest amount
Solute
substance present in lesser amounts
