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A&P I > 2. Inorganic Chemistry > Flashcards

2. Inorganic Chemistry Flashcards

Apply basic chemical concepts to the study of human physiology

1
Q

Define matter

A

has mass and takes up space

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2
Q

Define chemical element

A
  • substance made of only one atom type

- most basic form of matter

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3
Q

List major elements

A
  • oxygen (O)
  • carbon (C)
  • hydrogen (H)
  • nitrogen (N)
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4
Q

List lesser elements

A
  • calcium (Ca)
  • phosphorus (P)
  • potassium (K)
  • sulfur (S)
  • sodium (Na)
  • chlorine (Cl)
  • magnesium (Mg)
  • iron (Fe)
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5
Q

List trace elements

A

-aluminum, boron, chromium, cobalt, copper, fluorine, iodine, manganese, selenium, molybdenum, silicon, tin, vanadium, zinc

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6
Q

Define an atom

A
  • building block of all matter

- individual nucleus with protons or neutrons with electrons orbiting nucleus

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7
Q

Discuss the structure of an atom by describing the major subatomic particles (protons, neutrons, electrons) and their typical arrangement within the atom

A
  • nucleus has positively charged protons and neutrally charged neutrons
  • nucleus holds all of atomic mass (# of protons + neutrons) and is positively charged
  • electrons are negatively charged and on outside of nucleus
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8
Q

Proton

A
  • positively charged

- within nucleus of an atom

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9
Q

Neutron

A
  • have no electrical charge (neutral)

- within nucleus of atom

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10
Q

Electron

A
  • negatively charged

- exist in specific areas in rings around nucleus

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11
Q

Define isotope

A

-form of element with different number of neutrons (and therefore mass)

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12
Q

Difference between stable and radioactive isotopes

A
  • stable isotopes have same number of protons but different number of neutrons [ex. normal mass of carbon is 12 (6 protons, 6 neutrons), an isotope of carbon is carbon 14 still has 6 protons, but has 8 neutrons]
  • unstable isotopes have extra neutrons, nucleus becomes unstable and falls apart, releasing energy that is radioactive (in medicine use this for imaging)
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13
Q

Define and explain significance of atomic number

A
  • number of protons (+)
  • identifies atom or element
  • if this changes (# of protons) the identity of element/atom changes
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14
Q

Define and explain significance of atomic mass number

A

-number of protons and neutrons (+, 0)

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15
Q

Ion

A

-charged particles that are created by gaining or losing electrons

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16
Q

Cation

A
  • in ionic bonds, if ion loses electrons, the charge is positive
  • positively charged ion
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17
Q

Anion

A
  • in ionic bonds, if ion gains electrons, the charge is negative
  • Negatively charged ion
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18
Q

Molecule

A
  • covalent bonds create molecules
  • chemical particles composed of 2+ atoms united by chemical bond
  • atoms may be identical (ex nitrogen= N2) or different (glucose= C6H12O6)
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19
Q

Compound

A
  • molecules composed of 2+ elements by chemical bond

- ex CO2 (2 atoms of different elements)

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20
Q

Difference between molecular formula and structural formula

A
  • molecular formula identifies elements and how many atoms of each are present in a molecule
  • ex. C2H6O

-structural formula shows the difference between molecules by showing the location of atoms
H H
-ex. H-C-O-C-H
H H

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21
Q

Define isomer

A
  • molecules with identical molecular formula, but different arrangement of their atoms (same molecular formula, different structural formula)
  • ex. ethanol and ethyl ether have same molecular formula (C2H6O), but are not interchangeable
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22
Q

Define chemical bond and explain how octet rule can be used to predict how atoms will behave chemically

A
  • forces that attract molecules to one another and hold them together
  • octet rule says the valence shell wants to be full (have 8 electrons) and will give or take electrons accordingly to become stable
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23
Q

Define valence shell

A
  • outer most energy level of atom (valence electrons)
  • 1st level= 2 electrons max, 2nd level 8 electrons, 3rd level 16 electrons
  • determines which atoms will bond by octet rule
24
Q

Describe and give examples of ionic bonds

A
  • attraction of cation to anion
  • ions are charged particles that are created by gaining or losing electrons, [gaining electrons = negative (anion), losing electrons = positive (cation)]
  • in ionic bonds are bonds where electrons are donated from one atom or element to another (2 atoms collide and one of the nuclei has a stronger attractive force and pulls electrons away from the other. Electrons on valence level have highest energy and weakest attraction to nucleus so they are likely to be donated.
  • ex. sodium wants to give one and will give to chlorine to have 8 electrons. sodium loses electron so it becomes 1+ charge and becomes cation, chlorine gains electron becoming 1-, making it an anion. Electron donation sets up the bond and when it happens, opposite charges attract to each other and hold them together
25
Q

Describe and give examples of covalent bonds (including single, double, triple, polar, and non polar and polar covalent bonds)

A
  • form by the sharing of electrons; want to share to make 8; when elements come together neither nucleus can fully take away electrons from the other so electrons move around both nuclei
  • single covalent bonds share single pair of electrons (hydrogen H-H)
  • double covalent bonds share 2 pairs of electrons (oxygen O=O)(carbon dioxide O=C=0)
  • triple covalent bonds share 3 pairs of electrons (nitrogen N_=N)
  • non polar covalent bonds: strongest of chemical bonds when shared electrons spend equal time around each nucleus (don’t give a charge)
  • polar covalent bonds: when shared electrons spend more time orbiting one nucleus than another, they lend their charge to the region they spend more time in
26
Q

Distinguish between inorganic compounds and organic compounds

A
  • inorganic compounds lack carbon dioxide (except carbon dioxide, bicarbonate, and carbonic acid)
  • organic compounds are made of carbon
27
Q

Describe and give examples of covalent bonds (including single, double, triple, polar, and non polar and polar covalent bonds)

A
  • form by the sharing of electrons; want to share to make 8; when elements come together neither nucleus can fully take away electrons from the other so electrons move around both nuclei
  • single covalent bonds share single pair of electrons (hydrogen H-H)
  • double covalent bonds share 2 pairs of electrons (oxygen O=O)
  • triple covalent bonds share 3 pairs of electrons (nitrogen N_=N)
  • non polar covalent bonds: strongest of chemical bonds when shared electrons spend equal time around each nucleus (don’t give a charge)
  • polar covalent bonds: when shared electrons spend more time orbiting one nucleus than another, they lend their charge to the region they spend more time in
28
Q

Describe and give examples of hydrogen bonds

A
  • weak attraction between a slightly positive hydrogen atom in one molecule and a slightly negative oxygen or nitrogen atom in another
  • ex. water molecules are weakly attracted to each other by hydrogen bonds
  • hydrogen bonds also form between different regions of same molecule, especially in large molecules (ex. DNA, protein)
29
Q

Describe and give examples of ionic compounds

A
  • Ionic compounds have two different atoms ionically bonded

- ex. sodium chloride NaCl

30
Q

Explain the dissociation (ionization) of an ionic compound

A
  • ionic bonds are weak and easily dissociate (break up) in the presence of something more attractive, like water
  • ex. ionic bonds of NaCl break down easily as salt dissolves in water because both Na+ and Cl- are more attracted to water molecules than each other
31
Q

Describe and give examples of covalent compounds

A

Covalent compounds have different atoms covalently bonded

-ex. carbon dioxide CO2, glucose C6H12O6, DNA, protein

32
Q

Define chemical reactions and explain significance of the following terms to chemical reactions: reactant, products, metabolism

A
  • Chemical reactions are how these bonds form
  • Reactants: things you start with that come together to react and yield chemical products
  • Metabolism is the cells ability to perform these chemical reactions. Bonds are formed/ built up (anabolic) or broken down (catabolic) between atoms
33
Q

Define energy and define the difference between potential and kinetic energy

A
  • Potential energy is stored energy and not doing work all the time (like water behind a dam)
  • Kinetic energy is energy of motion or energy that is doing work (like water flowing out of a dam)
34
Q

Describe energy transfer in chemical reactions by explaining exergonic and endergonic reactions

A
  • exergonic reactions: energy releasing (catabolism)
  • ex. hold beaker of water and add sulfuric acid, beaker will get too hot
  • endergonic reactions: require energy input (anabolism)
  • ex. production of protein or fat
  • often coupled; energy released by exergonic is same energy absorbed in endergonic
35
Q

Describe energy transfer in chemical reactions by explaining activation energy and catalysts

A
  • activation energy: every reaction involves an activation energy to get started, so after initial cost end up releasing a lot on energy while others have initial cost and absorb even more beyond that
  • catalysts: substances that temporarily bind to reactants and hold them in a favorable position to react with each other; lower activation energy, allowing action to proceed faster because it’s done at a lower cost
36
Q

Describe energy transfer in chemical reactions by explaining law of conservation of energy

A

-the energy we put into a reaction should be equal to what we get out of it

37
Q

Describe different types of chemical reactions: synthesis reactions (anabolism)

A

-consists of energy storing synthesis reactions (require energy to input/ endergonic)

38
Q

Describe different types of chemical reactions: decomposition reactions (catabolism)

A

-consists of energy releasing decomposition reactions that break down covalent bonds, produce smaller molecules from larger ones, and release energy that can be used for physical work (exergonic)

39
Q

Describe different types of chemical reactions: exchange reactions

A

-2 molecules exchange atoms or groups of atoms

40
Q

Describe different types of chemical reactions: reversible reactions

A

-can go in either direction under different circumstances

41
Q

Distinguish between inorganic compounds and organic compounds

A
  • inorganic compounds lack carbon dioxide (except carbon dioxide, bicarbonate, and carbonic acid)
  • organic compounds are made of carbon
42
Q

Describe components of solution by explaining difference between solvent and solute

A
  • Solvent: substance doing the dissolving (ex. water)

- Solute: substance dissolved in solution (ex. sugar)

43
Q

Describe the components of the water molecule to the living system by describing the following characteristics of water: solvent, medium for chemical reactions, heat capacity, lubricant

A
  • solvency: ability to dissolve into other chemicals (water is universal solvent); biological molecules must be dissolved in water to move, come together, and react; body’s primary means of transporting substances
  • chemical reactivity of water is its ability to participate in chemical reactions; water ionizes into H+ and OH-, these can be incorporated into other molecules or released from them in the course of chemical reactions like dehydration synthesis
  • has a high heat capacity so it can absorb and release heat without changing temp; for temp to increase, molecules must move around, but hydrogen bond of water molecules inhibit their movement, so water can absorb heat without changing temp.
44
Q

Describe components of solution by explaining difference between solvent and solute

A
  • Solvent is substance doing the dissolving (ex. water)

- Solute is substance dissolved in solution (ex. salt)

45
Q

Define and give example of solution

A
  • mixtures that are uniform
  • solute and solvent are not visibly distinguished
  • don’t scatter light easily, usually transparent
  • can pass through most selectively permeable membranes
  • solute and solvent don’t separate when at rest
  • ex. water and sugar
46
Q

Define and give example of colloid

A
  • particles are a little larger but cannot be seen with naked eye
  • scatter light, usually cloudy
  • particles too large to pass through selectively permeable membrane
  • still small enough to remain mixed when at rest
  • ex. albumin in blood plasma, milk
47
Q

Define and give example of suspension

A
  • particles larger in size, visible to naked eye, make suspensions cloudy or opaque, too large to penetrate selectively permeable membrane, particles too heavy to remain suspended, so separate when standing
  • ex. blood cells in body
48
Q

Define and give examples of acids

A
  • any proton donor, a molecule that releases a proton (H+) in solution
  • amount of hydrogen ions in measured by pH scale (the more present/ higher the concentration, the lower the pH –> more acidic)
  • ex. hydrochloride HCl –> H+ + Cl-
49
Q

Define and give example of bases

A

-remove hydrogen ions from solution by releasing hydroxide ion (OH-) and those hydroxide ions will bind with hydrogen ions and take them out of solution, making H2O, neutralizing acid present
-ex. sodium hydroxide NaOH –> Na+ + OH-
makes base because H- can soak up any hydrogen ions to raise pH

50
Q

Define and give example of salts

A
  • dissociate into cations (+charged) and anions (-charged)

- ex. magnesium sulfate MgSO4 –> Mg2+SO42-

51
Q

Define pH and understand importance of hydrogen ion concentration to pH concept. Understand how pH scale can be used to describe solutions acidity or alkalinity.

A
  • pH is a measure derived from the molarity of H+
  • pH scale is 0-14; 7 is neutral
  • below 7 is acidic, the lower the more hydrogen ions
  • above 7 is basic (alkaline), the higher the less hydrogen ions
52
Q

Define buffer and explain the importance of buffer systems in maintaining homeostasis

A
  • buffers are chemical solutions that resist change in pH, lessening strength of acids and bases
  • convert strong acids or bases to weak ones by adding (make more basic) or removing (make more acidic) hydrogen ions
  • ex. small intestine needs to be basic enviro, stomach has a lot of acid so small intestine releases buffers to lessen acid by picking it up and neutralizing it
53
Q

Distinguish between proton donor and proton acceptor

A
  • proton donor:

- proton acceptor:

54
Q

Distinguish between hydrogen ion and hydroxide ion

A
  • hydrogen ion:
  • hydroxide ion:
  • hydrogen bonds are involved in attracting water molecules to hydroxide ions
55
Q

Define buffer and explain the importance of buffer systems in maintaining homeostasis

A
  • buffers are chemical solutions that resist change in pH, lessening strength of acids and bases
  • convert strong acids or bases to weak ones by adding (make more basic) or removing (make more acidic) hydrogen ions
  • ex. small intestine needs to be basic enviro, stomach has a lot of acid so small intestine releases buffers to lessen acid by picking it up and neutralizing it

A&P I (7 decks)

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