Chemistry of Life (Test 1) Flashcards
1
Q
Define matter.
A
Anything that takes up space and has mass.
2
Q
What is matter made up of?
A
Elements
3
Q
Define element.
A
A substance that cannot be broken down into other substances by chemical reactions.
4
Q
How many elements occur naturally in nature?
A
92
5
Q
Define compound.
A
A substance consisting of two or more different elements combined in a fixed ratio.
6
Q
Define emergent properties.
A
A compound has characteristics different from those of its elements.
7
Q
Explain the emergent properties of sodium chloride (NaCl).
A
Sodium is a metal that is highly reactive when exposed to water. Chlorine is a poisonous gas. When these two elements combine to form NaCl, they become edible table salt that has much different properties from the two elements it is composed of.
8
Q
Explain the difference in mass and weight.
A
Mass is the amount of matter in an object, whereas the weight of an object is how strongly that mass is pulled by gravity.
9
Q
Define essential elements.
A
Elements that an organism needs to live a healthy life and reproduce.
10
Q
What percent of the 92 naturally occurring elements are essential elements?
A
20-25%
11
Q
How many essential elements are there for humans? Plants?
A
Humans: 25
Plants: 17
12
Q
What elements make up 96% of living matter? What elements make up most of the remaining 4%?
A
96%: oxygen, carbon, hydrogen, nitrogen
4%: calcium, phosphorus, potassium, sulfur
13
Q
Define trace elements.
A
Elements required by an organism in only minute quantities.
14
Q
What trace element is needed by all forms of life?
A
Iron (Fe)
15
Q
In vertebrates, what trace element is an essential ingredient of a hormone produced by the thyroid gland?
A
Iodine (I)
16
Q
What does an iodine deficiency cause?
A
It causes the thyroid gland to grow to an abnormal size, a condition called goiter.
17
Q
In the human body, what are the 11 most abundant elements and their percentage of body mass?
A
O: 65%
C: 18.5%
H: 9.5%
N: 3.3%
Ca: 1.5%
P: 1.0%
K: 0.4%
S: 0.3%
Na: 0.2%
Cl: 0.2%
Mg: 0.1%
18
Q
Define atom.
A
The smallest unit of matter that still retains the properties of an element.
19
Q
What subatomic particles make up an atom?
A
neutrons, protons, and electrons
20
Q
What is the charge of a proton?
A
+
21
Q
What is the charge of a neutron?
A
no charge
22
Q
What is the charge of an electron?
A
-
23
Q
Define atom nucleus.
A
A dense core at the center of an atom where the protons and neutrons are packed together tightly.
24
Q
Where are the electrons located in an atom?
A
The electrons are rapidly moving around the nucleus where they form a “cloud” of negative charge.
25
What keeps the electrons in the vicinity of the nucleus?
The attraction between the opposite charges of the proton and electron.
26
What is the mass of a proton?
1.7 x 10^-24 gram (g) = 1 dalton = 1 amu
27
What is the mass of a neutron?
1.7 x 10^-24 gram (g) = 1 dalton = 1 amu
28
What is the mass of an electron?
About 1/2000 that of a proton or neutron
29
What does amu stand for?
atomic mass unit
30
What subatomic particle is unique to each element?
protons
31
Define atomic number.
(Z) The number of protons in an element. It can also indicate the number of electrons in an electrically neutral atom.
32
Define mass number.
(A) The sum of protons and neutrons in an atom's nucleus. This is also an approximation of the total mass of an atom (atomic mass).
33
Define isotopes.
Different atomic forms of the same element. All atoms of an element contain the same number of protons, but some atoms have more neutrons than other atoms of the same element and therefore have greater mass.
34
Define radioactive isotope.
An isotope in which the nucleus decays spontaneously, giving off particles and energy. When the radioactive decay leads to a change in the number of protons, it transforms the atom to an atom of a different element.
35
What does carbon-14 decay into?
nitrogen-14
36
Define half-life.
The time it takes for 50% of the parent isotope to decay.
37
Is half-life affected by environmental variables?
No, half-life is not affected by temperature, pressure, or any other environmental variables.
38
What subatomic particle is directly involved in chemical reactions?
electrons
39
Define energy.
The capacity to cause change by doing work.
40
Define potential energy.
The energy that matter possesses because of its location or structure.
41
What state of potential energy does matter move toward?
Matter has a natural tendency to move toward the lowest possible state of potential energy.
42
Why do electrons have potential energy?
Because of their distance from the nucleus. The further away from the nucleus, the more potential energy an electron possesses.
43
Define electron shells.
A place electrons are found, each with a characteristic average distance and energy level. An electron can only be found within this space, not between them.
44
How can electrons move between electron shells?
By absorbing or losing an amount of energy equal to the difference in potential energy between its position in the old shell and that in the new shell.
45
When an electron loses energy and moves to a lower electron shell, how is the energy released?
It is released to the environment as heat.
46
Define valence electrons.
The outermost electrons that determine most of the chemical behavior of an atom.
47
Define valence shell.
The outermost electron shell.
48
Max number of electrons in the first shell.
2 electrons
49
Max number of electrons in the second shell.
8 electrons
50
What is special about an atom with a completed valance shell?
It is unreactive or will not react readily with other atoms. The elements with a full valance shell are said to be inert, or chemically unreactive.
51
Define orbital.
The three-dimensional space where an electron is found 90% of the time.
52
How many electrons can occupy a single orbital?
2 electrons
53
Where does the reactivity of an atom arise from?
It arises from the presence of unpaired electrons in one or more orbitals of its valence shell.
54
Define chemical bonds.
An interaction in which atoms either share or transfer valence electrons which results in them staying close together.
55
Define covalent bond.
The sharing of a pair of valence electrons by two atoms.
56
Define molecule.
Two or more atoms held together by covalent bonds.
57
Define double bond.
When atoms form a molecule by sharing two pairs of valence electrons.
58
Define valence.
The bond capacity of an atom and usually equals the number of unpaired electrons required to complete the atom's valence shell.
59
Define electronegativity.
The attraction of a particular atom for the electrons of a covalent bond.
60
Define nonpolar covalent bond.
The electrons are shared equally because the two atoms have the same electronegativity.
61
Define polar covalent bond.
The electrons are not shared equally because the two atoms have different electronegativity. The amount of polarity depends on the relative electronegativity between the atoms that are bonded.
62
What does a polar covalent bond lead to?
It leads to partial negative charges on the more electronegative atom and a partial positive charge on the less electronegative atom.
63
Define ions.
Charged atoms or molecules.
64
Define cation.
Positively charged ion.
65
Define anion.
Negatively charged ion.
66
Define ionic bond.
The attraction between cations and anions due to their opposite charges. The transfer of the electron is not the bond, but rather it allows for a bond to form because of the ion formation. The ions do not need to have acquired their charge by an electron transfer with each other.
67
What are compounds formed by ionic bonds called?
ionic compounds or salts
68
How are salts found in nature?
As crystals of various shapes and sizes.
69
Can anything affect the strength of ionic bonds?
Yes, the environment plays a major factor in the strength of ionic bonds.
70
Define hydrogen bond.
When a hydrogen atom is covalently bonded to an electronegative atom, it becomes partially positive. This allows it to interact with electronegative elements in another molecule to form a weak bond.
71
Define van der Waals interactions.
Weak bonds that occur when atoms and molecules are very close together. Can be nonpolar atoms because, at any point in time, the electrons may cause split-second partial charges.
72
What does the cumulative effect of weak bonds reinforce?
The three-dimensional structure of molecules.
73
Why is molecular shape crucial to biological molecules?
It determines how biological molecules recognize and respond to one another with specificity. They often bind temporarily to each other by forming weak bonds, but only if their shapes are complementary.
74
Define chemical reactions.
The making and breaking of chemical bonds, leading to changes in the composition of matter.
75
Define reactants.
The starting materials in a chemical reaction.
76
Define products.
The final materials in a chemical reaction.
77
What affects the rate of a reaction?
The concentration of reactants.
78
Define chemical equilibrium.
The point at which the forward and reverse reactions offset one another exactly.
79
What is the only common substance to exist in the natural environment in all three physical states of matter?
Water
80
What is a polar molecule?
A molecule whose overall charge is unevenly distributed.
81
What are the four most important emergent properties of water?
1. cohesive behavior
2. ability to moderate temperature
3. expansion upon freezing
4. versatility as a solvent
82
What is cohesion?
The attachment of water molecules to other water molecules. As water evaporates from a leaf, hydrogen bonds cause water molecules leaving the veins to tug on molecules further down, and the upward pull is transmitted through the water-conducting cells all the way to the roots.
83
What is adhesion?
The clinging of one substance to another. Adhesion by hydrogen bonds to the molecules of cell walls helps counter the downward pull of gravity.
84
What is surface tension?
A measure of how difficult it is to stretch or break the surface of a liquid. Related to cohesion.
84
How does water moderate air temperatures?
By absorbing heat from air that is warmer and releasing the stored heat to air that is cooler. It is effective as this because it can absorb or release a large amount of heat with only a slight change in its own temperature.
85
What is kinetic energy?
The energy of motion. The faster a molecule moves, the greater its kinetic energy.
85
What is thermal energy?
The kinetic energy associated with the random movement of atoms or molecules.
86
What is temperature?
A measure of energy that represents the average kinetic energy of the molecules in a body of matter.
87
How does thermal energy pass from object to object?
Thermal energy passes from the warmer to the cooler object until the two are the same temperature. The thermal energy in this transfer is called heat.
88
What is a calorie (cal)?
The amount of heat it takes to raise the temperature of 1g of water by 1 C, or the amount of heat that 1g of water releases when it cools by 1 C.
89
What does 1 J equal?
0.239 cal;
1 cal = 4.184 J
90
What is specific heat?
The amount of heat that must be absorbed or lost for 1g of that substance to change its temperature by 1 C.
91
What is the specific heat of water?
1 cal/g*C
92
Why is a high specific heat of water important to humans?
Because humans are made mostly of water, they are better able to resist changes in their own temperature.
93
What is evaporation?
The transformation from a liquid to a gas.
94
What is heat of vaporization?
The quantity of heat a liquid must absorb for 1g of it to be converted from the liquid to the gaseous state.
95
How much heat is needed to evaporate 1g of water at 25 C?
580 cal
96
What is evaporative cooling?
As a liquid evaporates, the surface of the liquid that remains behind cools down. This is because the molecules with the greatest kinetic energy are the most likely to leave as a gas.
97
What happens to ice in liquid water?
It floats because water expands when it solidifies. At 0 C, it locks into a crystalline lattice, with each hydrogen bonded to four partners.
98
What is a solution?
A liquid that is a completely homogeneous mixture of two or more substances.
99
What is a solvent?
The dissolving agent of a solution.
100
What is a solute?
The substance that is dissolved in a solution.
101
What is an aqueous solution?
A solution in which the solute is dissolved in water.
102
Why is water a very versatile solvent?
Because of the polarity of the water molecules.
103
What is a hydration shell?
The sphere of water molecules around each dissolved ion.
104
What does it mean to be hydrophilic?
Any substance that has an affinity for water.
105
What does it mean to be hydrophobic?
Substances that are nonionic and nonpolar that repel water.
106
What is molarity (M)?
The number of moles of solute per liter of solution.
107
What is an acid?
A substance that increases the hydrogen ion concentration of a solution.
108
What is a base?
A substance that reduces the hydrogen ion concentration of a solution.
109
Formula for pH.
pH = -log [H+]
110
What is a buffer?
A substance that minimizes changes in the concentrations of H+ and OH- in a solution.
