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HESI A2 > Chemistry > Flashcards

Chemistry Flashcards

1
Q

What is chemistry?

A

The study of matter and its properties, and how those properties relate to one another.

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2
Q

Everything in the universe is made or composed of different kinds of matter in one of its three states:

A

liquid, solid, gas.

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3
Q

What is matter defined by?

A

Its properties

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4
Q

—— is the scientific system of writing numbers- very big or very small- easily.

A

Scientific notation

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5
Q

Scientific notation is composed of three parts:

A
  1. a mathematical sign (+ or -)
  2. significand
  3. exponential/ logarithm
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6
Q

What is the significand?

A

The base value of the number or the value of the number when all the values of ten are removed.

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7
Q

What is the exponential?

A

A multiplier of significand in powers of ten.

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8
Q

Difference between a positive exponential and a negative exponential?

A

Positive multiplies the significand by factors of ten.
Negative multiplies the significand by factors of one tenth (0.1).

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9
Q

HINT for exponents representing exponentials:

A

The exponent shows how many 0s are present in the value. For negative exponents, the 0 before the decimal counts as one of the zeros present in the value.
EX: 10^6 = 1,000,000
10^ -6 = 0.000001

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10
Q

Consider the following:
-9.0462 X 10^5 =
What is the significand and what is the final answer?

A

The significand is -9.0462
The final answer is -904,620

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11
Q

Consider the following:
4.7 x 10^-3
What is the final value?

A

0.0047

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12
Q

Exponent hint when multiplying the significand by a positive vs a negative exponent:

A
  • exponent shows how many places to move the decimal to the left
    + exponent show how many places to move the decimal the the right.
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13
Q

What system measure weight, length, and volume?

A

Metric system. Standard in health professions

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14
Q

3 basic measurements of the metric system:

A

-grams (weight)
-liters (volume)
-meters (distance)

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15
Q

Label the prefix and the basic unit of measure in “kilogram”

A

Prefix: kilo
basic unit of measure: gram

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16
Q

What are prefixes?

A

Quantifiers of the measurement units.

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17
Q

What prefix represents 10^12, or 1 trillion times

A

Tera (T)

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18
Q

What prefix represents 10^9; 1 billion times?

A

Giga (G)

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19
Q

What prefix is for 10^6; 1 million times?

A

Mega (M)

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20
Q

What prefix is for 10^3; 1 thousand times?

A

kilo (k)

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21
Q

What prefix if for 10^2; 100 times?

A

hecto (h)

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22
Q

What prefix is for 10^1; ten times?

A

Deka (D)

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23
Q

What prefix is for 10^-1; 1 tenth of

A

deci (c)

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24
Q

What prefix is for 10^-2; 1 hundredth of

A

centi (c)

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25
Q

What prefix is for 10^-3; 1 thousandth of?

A

milli (m)

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26
Q

What prefix is for 10^-6; 1 millionth of?

A

micro (u)

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27
Q

What prefix is 10^-9; 1 billionth of?

A

nano (n)

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28
Q

What prefix is for 10^-12; 1 trillionth of?

A

pico (p)

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29
Q

What prefix is for 10^-15; 1 quadrillionth of?

A

femto (f)

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30
Q

What is a little less than 2 cm in diameter?

A

A dime

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31
Q

About how many pounds are in a kilogram?

A

2.2 lbs

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32
Q

A liter is a little more than a —-

A

quart

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33
Q

Three most common temperature systems:

A

Fahrenheit, Celsius, Kelvin

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34
Q

Zero degrees F is what?

A

The freezing point of sea water or heavy brine at sea level

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35
Q

What is the freezing point of pure water at sea level in Fahrenheit?

A

32 degreesF (0 C) (273K)

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36
Q

What is the boiling point of pure water at sea level in Farenheit?

A

212 degrees F (100 C) (373K)

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37
Q

What is the normal body temperature in Fahrenheit?

A

98.6 degrees F (37 C) (310K)

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38
Q

Celsius is sometimes called?

A

Centigrade: used by the rest of the world and the scientific community.

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39
Q

What temperatures are thought to be the lowest temperature achievable? or “absolute zero” (o)?

A

-273.15 C ; 0K

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40
Q

What is the basic building block of all molecules?

A

An atom

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41
Q

Physical structure of an atom?

A

Nucleus, orbits, and sometimes electron clouds.

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42
Q

Electrons orbit the nucleus at various energy levels called:

A

shells or orbits (like layers of an onion)

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43
Q

When is an atom considered to be most stable?

A

When its outermost orbit is full.

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44
Q

What type of atoms tend to have equal numbers of protons and electrons, making them electrically neutral?

A

“Ground state” atoms

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45
Q

What is an atom called when it is electrically charged?

A

An ion; is in an ionic state

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46
Q

Words indicating charges of an atom as positive or negative.

A

Negative : “anion”
Positive: “cation”

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47
Q

The rows of a periodic table are called:

A

Periods: horizontal.
They have similar physical properties and the same number of shells.
7 periods

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48
Q

The columns in the periodic table are called:

A

Groups: vertical.
Reacts to other elements in similar ways.
18 groups

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49
Q

What is the periodic table?

A

A chart of known elements arranged according to their properties; helps to predict charge of an atom or element when it exists as an ion.

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50
Q

How do the groups progress?

A

IA (+1); IIA (+2); IIIA (+3);
IVA (+4) OR (-4);
VA (-3); VIA (-2); VIIA (-1);
VIII = Noble gases (no charge) remains neutral in all solutions/situations.

The chart also shows how many electrons are likely to be in the outer shells of each by the group.

Rows show how many shells they have.

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51
Q

Groups 3 IIIB through 12 IIB on the periodic chart are what?

A

Transition metals: not as straightforward to predict because of some exceptions to the rules.

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52
Q

What are two important numbers or properties of atoms that can be obtained from the periodic table?

A

Atomic number
Atomic mass

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53
Q

What is the atomic number?

A

The number of protons in the nucleus. Defines an atom as a particular element.
EX: Any element with 11 protons, regardless of neutrons or electrons, is sodium (Na).

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54
Q

Where is the atomic number located?

A

At the top of each square in the periodic table. (Always a whole number).

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55
Q

What is the atomic mass?

A

The average mass of each of that element’s isotopes.

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56
Q

What is an isotope?

A

Different kinds of the same atom that vary in weight due to varying numbers of neutrons (always the same number of protons).
EX: Carbon 12 (^12 C), the most common carbon has 6 protons and 6 neutrons.
The isotope used for carbon dating is Carbon 14 (^14 C), which has 6 protons and 8 neutrons.

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57
Q

What makes up the elements mass?

A

Protons and neutrons makes up nearly all of the atom’s total mass.

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58
Q

Where can the atomic mass be found?

A

At the bottom of each of the squares. (Is usually a decimal).

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59
Q

What is usually seen of protons and neutrons with elements found in their most common form?

A

The number of neutrons and protons is the same.

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60
Q

What is the simplest form of matter that can naturally exist in nature?

A

An element or an atom.

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61
Q

When elements/atoms exist in combination with other elements, they are called:

A

Compounds: They combine in whole number ratios.

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62
Q

Chemical equations are made up of two parts:

A

Reactants: react to produce desired end results or compounds
Products: The end results or compounds
Reactants&raquo_space;> Products

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63
Q

What law states that mass cannot be created or destroyed during a chemical reaction?

A

The law of conservation of mass. This means that a chemical equation must be balanced on both sides of the equation.

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64
Q

What is an example of equilibrium?

A

A chemical reaction that proceeds to completion, leaving some reactants without products.

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65
Q

In equilibrium, what is the rate in which reactants are forming products, and products are forming reactants?

A

The same rate.

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66
Q

A reaction at equilibrium is said to be what?

A

Reversible.

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67
Q

Through manipulation of the reaction by various mean, shifts in equilibrium reversibility or the rate of the reaction can be controlled. What are the 4 basic ways to increase reaction rate?

A

increase temperature in the reaction
increase surface area of the reactants
add a catalyst
increase concentrations of the reactants

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68
Q

How does increasing the temperature increase the reaction rate?

A

Increase in temperature = particles have greater kinetic energy.
Moving faster = greater chance of collision at higher energy forces.

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69
Q

When do chemical reactions occur?

A

Upon contact

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70
Q

How does increasing the surface area increase reaction rate?

A

Smaller particles of reactants allows for greater chances of contact.

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71
Q

How does a catalyst increase reaction rate?

A

Reduces activation energy/ amount of energy necessary for a reaction to occur.

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72
Q

What happens to a catalyst after a reaction?

A

It is not fully used up and can be collected at completion.

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73
Q

Examples of common catalysts:

A

metals and proteins (enzymes)

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74
Q

How does increasing the concentration of the reactants increase reaction rate?

A

More of a particle will add to the likelihood of contact/collision between reactions, producing more products.

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75
Q

—— is a homogenous mixture of two or more substances.

A

Solution

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76
Q

What makes up a solution?

A

Solute and solvent

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77
Q

——- is the part/parts of a solution that is being dissolved.

A

Solute

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78
Q

—— is the part of a solution that is doing the dissolving

A

Solvent

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79
Q

solutions can be made up of what matter forms?

A

liquid in a liquid
solid in a liquid
solid in a solid

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80
Q

4 types of solutions:

A

Compounds
Alloys
Amalgams
Emulsions

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81
Q

—– is a type of solution made of mixtures of different elements to create a single matter

A

Compounds

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82
Q

—— is a type of solutions that is solid and made up of metals to make a new one such as bronze (copper and tin), or steel (iron and carbon)

A

Alloys

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83
Q

Alloys may contain the following metals: (3)

A

Tungsten, chromium, and manganese

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84
Q

—– is a solution that is a specific type of allow in which a metal is dissolved in mercury.

A

Amalgams

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85
Q

—– is a solution type that is a mixture of matter that readily separates such as water and oil

A

Emulsions

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86
Q

—— is expressed as weight per weight (e.g. grams per grams); weight per volume (e.g. grams per liters); or volume per volume (ml per L)

A

Concentration

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87
Q

—— ——– is the expression of concentrations as parts per 100 parts.

A

Percent concentration
EX: mg per 100 mL = mg/100mL =mg/dL

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88
Q

A concentration expression of milliliters (mL) per 100 milliliters (mL) can be written as?

A

mL/100mL or mL/dl

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89
Q

What is one of the most important concepts in chemistry?

A

A mole

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90
Q

What is a more sophisticated way to express concentrations than percent?

A

Molarity, or molar concentration

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91
Q

What is a mole?

A

6.02 x 10^23 molecules of something. More than a trillion trillions
“Avogadro’s number”

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92
Q

A one molar solution will contain 6.02 x 10^23 representative molecules of a solute in how much of a solvent?

A

1 liter

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93
Q

How are molar concentrations written?

A

mol/L

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94
Q

one mole, or 6.02 x 10^23, atoms of an element or compound is equal to what?

A

The atomic mass of the element/compound in grams = 1 mole

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95
Q

——– is the making or changing of chemical bonds between elements/compounds to create new chemical compounds- resulting in different chemical formulas and different chemical properties.

A

Chemical reaction.

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96
Q

5 main types of chemical reactions:

A

Synthesis, decomposition, combustion, single replacement, double replacements.

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97
Q

When a reaction occurs, what is the product generally?

A

A molecule.

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98
Q

A molecule in a chemical equation may have what after the chemical symbol?

A

A subscript (opposite of exponent position on the right hand side)
EX: O(2) = oxygen

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99
Q

What type of reaction represents two elements combining to for a product?

A

Synthesis:
EX: Potassium and chloride combine to make potassium chloride salt
2K+ + 2Cl-&raquo_space; 2KCl (notice the + and the - works to make a whole molecule. The common value of both molecules can be placed as a multiple in front of the whole product)

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100
Q

——- is a type of chemical reaction that is the opposite of synthesis because it breaks a compound into its component parts.

A

Decomposition:
EX: NaCl&raquo_space;> Na+ + Cl-
(cation and anion)

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101
Q

This type of chemical reaction is self-sustaining, exothermic (creates heat) reaction where oxygen and a fuel compound such as hydrocarbon react.

A

Combusiontion.

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102
Q

In the combustion of hydrocarbon (gas or oil product), what are the productes?

A

CO2 and H2O (ALWAYS)
CH + O2 > CO2 + H2O (with stoichiometry to determine quantities of each to balance the equation)

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103
Q

What does the combustion equation of ethane look like?
Ethane = 2C(2)H(6) (g) + 7O(2) (g)
Combustions must react with oxygen. (g) = gas

A

Start with Carbon.. 2C(2) = 4 carbon available.
We know carbon dioxide makes CO2. How many oxygen is available?
7O(2) = 14 oxygen molecules available.
For every carbon molecule, we need to oxygen molecules. : 4C + 8O: = 4CO(2) (g)
How many oxygen molecules are left? 6O
How many hydrogen molecules are in ethane? 12. 2H(6) = 12H
For H(2)O we end up with 6H(2)O (g)

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104
Q

Type of chemical reaction that involves ionic compounds.

A

Single replacement

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105
Q

What determine whether or not a reaction will take place in a single replacement reaction?

A

The reactivity of the metals involved.

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106
Q

A single replacement reaction consists of a more active metal that reacts with what to produce a new compound?

A

An ionic compound that made up of a less active metal.

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107
Q

What is a good example of a single replacement reaction?

A

Combining copper (Cu) with aqueous silver nitrate (AgNO(3)).

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108
Q

What will happen when combining copper (Cu) with aqueous silver nitrate (AgNO(3)).

A

The copper and the nitrogen will switch places.
(aq) stands for aqueous
(s) stands for solid

Cu (s) + 2AgNO(3) (aq)&raquo_space; Cu(NO(3)(2)) (aq) + 2Ag (s)

Copper + silver nitrate yields copper nitrate + silver

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109
Q

Type of reaction that involves two ionic compounds that have “switched partners” where the positive ion from one compound combines with the negative ion of the other compound (and vice-versa).

A

Double replacement

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110
Q

Good example of a double replacement chemical reaction?

A

Silver nitrate + potassium chloride
yields silver chloride + potassium nitrate
AgNO(3) + KCl&raquo_space;> AgCl + KNO(3)

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111
Q

—— is the joining of one atom, element, or chemical to another

A

Chemical bonding.

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112
Q

The type of bonding that occurs is determined by what?

A

The electrons in the outer shell of the atom.

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113
Q

Two main types of chemical bonds:

A

ionic and covalent

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114
Q

——- is a type of bond that is an electrostatic attraction between two oppositely charged ion (cation and anion). “Taking and giving”

A

ionic bond. also known as an “electrostatic bond”
The gain and loss of an electron to make up a neutral charge

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115
Q

Ionic bonds are usually formed between what two elements?

A

A metal and a nonmetal

116
Q

What a good example of an ionic bond?

A

Salt: Na+ + Cl- = NaCl
The + means that Na is missing 1 electron from its outer shell.
The - means that Cl has one electron on its outer shell. Cl give an electron to Na making them both stable with their outer shells being completely inhabited through their bond.

117
Q

What type of bond is formed when two atoms “share” electrons. generally in pairs, with one pair from each atom?

A

Covalent bond

118
Q

Types of covalent bonds:

A

Single, double, and triple:
Single shares one pair
Double shares two pairs
Triples shares three pairs of electrons.

119
Q

What is the strength of a covalent bond and what elements usually form a covalent bond?

A

Strongest of any chemical bond.
Formed between two nonmetals.

120
Q

What determines a covalent bond to be non-polar?

A

If the electrons within the bonded compound are shared equally.

121
Q

What determines a polar covalent bond?

A

The shared electron density of the bond is concentrated around one atom more than the other. This gives the compound different charges on different sides. (more - on one side; more positive on another)

122
Q

What is polarity based on?

A

Difference in electronegativity values amongst the elements involved in the bond. The greater the difference = the more polar the bond will be.

123
Q

Example of a nonpolar covalent bond.

A

H2. Both are missing one electron from their outer shell. By combining to for H(2), their share one an electron with each other (single covalent bond), fulfilling their outer shell.
H-H

124
Q

Give an example of a polar covalent bond.

A

CO2 or O=C=O.

Carbon has only 4 electrons on it’s outer shell (6 electrons total). 2 Oxygen molecules each have 6 electrons in their outer shell (8 electrons total). Each Oxygen molecule shares one pair of electrons with each carbon, while each carbon shares one pair with one oxygen. This forms a double bond.

125
Q

—– are weaker forces of attraction between whole molecules (not bonds).

A

Intermolecular foces

126
Q

Three types of intermolecular forces:

A

hydrogen bonds, dipole-dipole interactions and dispersion forces

127
Q

This type of molecular force is the attraction for a hydrogen atom by a highly electronegative element

A

Hydrogen bond

128
Q

What elements are generally involved in a hydrogen bond?

A

Fluorine (F), oxygen (O), and nitrogen (N)

129
Q

How strong are hydrogen bonds compared to covalent bonds?

A

5%-10% as strong as covalent bonds (strongest of the intermolecular forces)

130
Q

A ——- is created when an electron pair is shared unequally in a covalent bond between two atoms or elements (polar covalent bonding).

A

A dipole

131
Q

——— is an intermolecular force made up of an attraction of one dipole on one molecule for the dipole of another molecule.

A

Dipole-Dipole Interactions

132
Q

What is the strength of a dipole-dipole interaction?

A

The negative end of a polar bonded molecule with a covalent bond is attracted to the positive end of a polar bonded molecule with a covalent bond. The attraction is a weak intermolecular force which is only about 1% as strong as a normal covalent bond.

133
Q

—– is an type of intermolecular force that is the weakest of all. It occurs when electrons within an element or compound concentrates themselves on one side of an atom creating a momentary/temporary dipole, which then is attracted to another momentary dipole of opposite charge in another nearby element/compound

A

Dispersion forces/ “London dispersion forces”

134
Q

——- is the part of chemistry that deals with the quantities and numeric relationships of the participants in a chemical reaction.
It is the calculation of relative quantities of reactants and products in a balanced chemical equation.

A

Stoichiometry

135
Q

Stoichiometry is similar to comparing ingredients in a recipe. For a chemical equation to be balanced, what is placed in front of each compound?

A

Numbers called coefficients: used in a ratio to compare how much of one substance is needed to react with another in a certain reaction.

136
Q

What process determines the number of moles of a certain element/compound that is needed to complete a reaction, thereby determining the coefficient for each (stoichiometry)

A

Dimensional analysis

137
Q

Using dimensional analysis, use this reaction to determine the number of moles of oxygen (O(2)) that will reach with (4) molecules of ethane (C2H6)
2C(2)H(6) + 7O(2)&raquo_space;> 4CO(2) + 6H2O
Ethane + oxygen yields carbon dioxide + water

A

2 x ethane must be done to oxygen. 14

138
Q

——– involves the transfer of electrons from one element to another.

A

Oxidation/reduction reactions/ “redox”

139
Q

—— is the loss of electrons

A

Oxidation

140
Q

—– is the gain of electrons

A

Reduction.

141
Q

It is impossible to have oxidation without reduction, and vice-versa. The element that is oxidized (loses electron) is the ——-, and the element that is reduced (gains electron) is the ——-

A

oxidized = reductant or reducing agent
reduced = oxidant or oxidizing agent

EX:
Oxidant (gains electron) + e- (is reduced) <-> Reductant (loses electron) - e- (is oxidized)

142
Q

To identify what has been oxidized and what has been reduced, the oxidation states of all elements in the compound must be determined. There are 5 rules that aids in making these determinations:
1. If an element is in its natural state, then its charge is —–.
Most elements in their standard states are single atoms with 6 exceptions. what are they?

A

Natural state: oxidation number/charge is zero.
6 elements that are exceptions to the existence commonly in it’s natural states is hydrogen (H2), bromine (Br2), oxygen (O2), nitrogen (N2), iodine (I2) and fluorine (F2).
When they exist outside of a compound in their natural state, they are always in pairs.

143
Q

Rule #2 to determine oxidation states is that the oxidation number of any simple ion is what?

A

The charge of the ion. EX: sodium (Na) is listed as an ion (Na+) = oxidation number of +1.
Cl has an oxidation number of -1

144
Q

The third rule for determining the oxidation number for oxygen in a compound is:

A

-2

145
Q

The oxidation number for hydrogen in a compound is:

A

+1

146
Q

The sum of the oxidation numbers equals what?

A

The charge on the molecules or polyatomic ions

147
Q

Assign oxidation numbers to all elements in the following reaction:
2C(2)H(6) + 7O(2) > 4CO(2) + 6H(2)O

A

*Ignore the coefficient. Has nothing to do with oxygen number
Rule 5 says that compounds (which likely do not have a charge) will have the same oxidation number as the compound. In this case ethane has a charge of 0.
Hydrogen oxidation is +1. In this equation, there are 6 hydrogens. hydration has an oxidation number of +6
To solve for carbon:
2(C) + 6(+1) = 0. Solve for C: -6/2 = -3
Oxygen is zero (rule one.
Carbon in CO(2) is +4 and O is -2.
Hydrogen is +1 in water and Oxygen is -2.
I do not understand.
The idea is that any atom that has changed charge has either had to gain or lose electrons. if - to + electrons are lost. if neutral or positive to - electrons are gained (reduced)

148
Q

—– are corrosive metals that change blue litmus paper red.

A

Acids

149
Q

——- are substances that denature proteins, making them feel very slick; they change red litmus paper to blue

A

Bases (alkaline compounds)

150
Q

Acids are compounds that donates what?

A

Protons or hydrogen

151
Q

Hydrogen in its ionic state is what?

A

a proton

152
Q

What happens to naked protons in water?

A

Naked protons exist only for a short time before reacting with other water molecules to produce H3O+ (hydronium).

153
Q

What is the makeup of hydronium?

A

A water molecule plus a proton or hydrogen

154
Q

What do bases do with hydrogen or protons and what type of group do they usually have?

A

Bases are hydrogen or proton acceptors
They generally have a hydroxide group.

155
Q

Explain the dissociation of water into low concentrations of hydronium and hydroxide ions:
water + water yields acid + base
H2O +H2O <-> H(3)O+ + OH-

A

One water molecule acts as a hydrogen donor, giving one of its hydrogens to another water molecule, which produces hydronium (H(3)O+) and leaves a hydroxyl group (OH-)

156
Q

What do all acids do when placed in water?

A

Produce hydronium.

157
Q

What is the word that describes how water can act as an acid and a base?

A

Amphoteric

158
Q

What is the goal of chemical reactions

A

To reach stable electron configurations.

159
Q

—– is concerned with reactions that take place in the nucleus and obtain stable nuclear configurations

A

Nuclear chemistry

160
Q

——- is the word used to describe the emission of particles and/or energy from an unstable nucleus.

A

Radioactivity

161
Q

The particles and/or energy that are emitted are referred to as what?

A

Radiation

162
Q

Three types of radiation in nuclear chemistry

A

Alpha, beta, gamma

163
Q

—– is the emission of helium nuclei.

A

Alpha radiation

164
Q

emission of helium nuclei are alpha particles that contain how many protons and neutrons?

A

two protons and two neutrons (+2).

165
Q

Of all the radioactive emissions, alpha is what?

A

The largest; penetration from alpha particles can generally be stopped by a piece of paper.

166
Q

——- is a product of decomposition of a neutron or proton.

A

Beta radiation

167
Q

Beta radiation is composed of what?

A

High-energy, high speed electrons that began as neutrons or potons.

168
Q

What is the charge of beta particles?

A

negatively charged (electrons) or positively charged (positrons)

169
Q

Why can beta particles be stopped by a thin sheet of aluminum foil, Lucite, or plastic?

A

No mass

170
Q

—- is high-energy electromagnetic radiation similar to x-rays but with more energy

A

Gamma radiation

171
Q

What is the penetration strength of Gamma rays?

A

Can go through several feet of concrete or several inches of lead. Lead shielding is required to block gamma rays.

172
Q

How are isotope written?

A

As an abbreviation with the symbol of the element preceded by a superscript number indicating atomic mass
EX Iodine-131 = ^131 I

173
Q

In nature, what is the stability of isotopes?

A

Some are stable, some are not.

174
Q

What happens to an unstable nuclei?

A

Will change/ “decay” into a more stable form.

175
Q

—– is the amount of time it takes for half of the unstable isotope to decay

A

“half-life”

176
Q

In nuclear chemistry the unstable atom decays until it finds a stable nuclear configuration, usually by doing what?

A

emitting radioactive particles

177
Q

What is the symbol for half life?

A

T^1/2

178
Q

What is the half-life of radioactive elements?

A

Different for each one. From very long to as short as a few days.

179
Q

If an isotope such as ^131 I has a half-life of approximately 8 days. what well occur every 8 days?

A

one-half of the radioactive particles will be emitted or decayed. Happening over and over until reaching a stable nuclear configuration.

180
Q

——- is the study of chemical processes in living organisms

A

Biochemistry

181
Q

Biochemistry deals with structures and functions of molecules such as what?

A

carbohydrates, proteins, lipids, and nucleic acids

182
Q

Sugar in DNA (deoxyribonucleic acid)

A

Deoxyribose

183
Q

Sugar in RNA (ribonucleic acid)

A

Ribose

184
Q

—– are more abundant than any other known type of biomolecule

A

Carbohydrates

185
Q

What is the simplest type of carbohydrate?

A

Monosaccarides

186
Q

Monosaccharides contain carbon, hydrogen and oxygen in what ration?

A

1:2:1 C6H12O6

187
Q

Two examples of monosacharades?

A

Glucose and Fructose- both have the c6H12O6 chemical formula, however have different molecular configurations-
. and are both six-carbon sugars called a hexose

188
Q

Where does the word saccharide come from?

A

The Greek word meaning “sugar”

189
Q

—— two monosaccharides joined together

A

Disaccharides

190
Q

Example of disaccharide:

A

Sucrose (table sugar) = glucose + fructose
Lactose: glucose + galactose
Maltose: Glucose + glucose

191
Q

When three to six monosaccharides are joined together it is a :

A

oligosaccharide. (oligo- means few)

192
Q

More than six and up to thousands of monsaccharides joined together make a

A

polysaccharide “starch”

193
Q

Two most common polysaccharides:

A

Cellulose (made by plants)
Glycogen (made by animals)
Both are chains of repeating glucose units

194
Q

Glucose is mainly metabolized by a chemical pathway in the body called?

A

Glycolysis

195
Q

What is the net result of the breakdown of one molecule of glucose in glycolysis?

A

Two pyruvate; Two adenosine triphosphate (ATP)

196
Q

If oxygen is available after glycolysis, the pyruvate is futher metabolized by a process called

A

oxidative phosphorylation (Krebs cycle)

197
Q

What molecules does the Kreps cycle produce?

A

ATP, water, carbon dioxide.

198
Q

When skeletal muscles are used in vigorous exercise, they will not have enough oxygen to meet their energy demands. They will need to use another type of glucose metabolism called:

A

anaerobic glycolysis.

199
Q

What occurs in anaerobic glycolysis?

A

Glucose is converted to lactate instead of pyruvate = creates “burning or cramping” sensation during intense exercise.

200
Q

What is a name for an anaerobic organism that can function in low concentrations of oxygen?

A

micro-aerobic

201
Q

The liver can make glucose from other noncarbohydrate sources, such as proteins and parts of fats, using a process called?

A

Gluconeogenesis

202
Q

What happens to glucose after gluconeogenesis?

A

Either used by cells to create ATP or can be stored as glycogen in animals or cellulose in plants.

203
Q

Glucose can be used to make what?

A

Other saccharides

204
Q

What is an amino acid composed of?

A

a carbon atom bonded with four other groups.

205
Q

An amino acid is a carbon atom bonded with four other groups which are:

A

amine group (NH(2))
carboxyl group (COOH)
hydrogen
R group

206
Q

What makes each amino acid different?

A

The R group

207
Q

A union of two amino acids using a peptide bond is called a :

A

dipeptide

208
Q

groups of fewer than 30 amino acids are called:

A

peptides or polypeptides

209
Q

Larger groups of amino acids (30 to thousands) is a

A

protein.

210
Q

An important protein in blood is albumin which contains how many amino acids?

A

585

211
Q

Fatty acids consist of what?

A

Hydrocarbon chain with an acid group (carboxyl group (COOH)) at one end.

212
Q

What is a neutral fat?

A

Triglyceride: 3 fatty acids joined to a glycerol or some other backbone structure

213
Q

Phospholipids are similar to neutral fats. What is different?

A

One of the three fatty acids is replaced by a phosphate group.

214
Q

What is a form of fat composed of a four-ring structure and a side chain?

A

Cholesterol

215
Q

5 functions of fat:

A

Insulate body organs against shock
Maintain body temperature
Keep skin and hair healthy
Promote cell function
Energy storage

216
Q

Examples of natural lipid classifications:

A

Unsaturated, polyunsaturated, saturated.

217
Q

—— have no double bond between carbon atoms of the fatty acid chains

A

Saturated fats

218
Q

—— have one or more double bonds between some of the carbon atoms of the fatty chains

A

Unsaturated fats

219
Q

—– are the biologic brain of life, telling the cell what it will do and how to do it.

A

Nucleic acids.

220
Q

Examples of nucleic acid structures:

A

RNA and DNA

221
Q

DNA & RNA are made up of what that conveys genetic information?

A

Nucleotide chains

222
Q

What is found in all living cells and viruses?

A

Nucleic acids

223
Q

Where can nucleic acids be found?

A

Most in the nucleus, some in the cytoplasm and mitochondria of individual cells.

224
Q

Nucleic acids are very large molecules that have two main parts. In DNA the composition is what?

A

Backbone: deoxyribose (five-carbon sugar called a pentose) and a phosphate. Alternately chains together making two long structures. (covalently bonded)
Base pairs (hydrogen bonded)

225
Q

RNA composition

A

Single strand of ribose (five-carbon carbohydrate), in a sugar-phosphate chain.

226
Q

What would be the oxidation state of the sulfur atom in sulfuric acid?
H(2)SO(4)?
A. +4
B. +5
C. +6
D. +8

A

+6 (do not forget to consider the ring structures. Sulfur has 16 protons. 2 in the first shell, 8 in the 2nd shell, which leaves 6 in the outer shell.

227
Q

The nucleic acids DNA and RNA:
A. Are found in the cell nucleus
B. Are not found in the cell mitochondria
C. Contain different kinds of fat
D. Include very small molecules

A

A. Are found in the cell nucleus

228
Q

Glycogen is:
A. A plant starch
B. A disaccharide
C. An animal starch
D. Not a starch or a disaccaride

A

C. An animal starch

229
Q

Which of the biologic molecules are considered the most significant contributors to cellular function?
A. Carbohydrates
B. Lipids
C. Proteins
D. Nucleic acids

A

C. Proteins

230
Q

Beta radiation is the emission of which of the following?
A. Large numbers of helium ions
B. An electron
C. High-energy electromagnetic radiation
D. A product of the decomposition of a proton

A

B. An electron

231
Q

What are most of the elements (besides hydrogen (H)) in the first group?

A

Alkali metals:
Li: lithium (3)
Na: Sodium (11)
K: Potassium (19)
Rb: Rubidium (37)
Cs: Caesium (55)
Fr: Francium (87)

232
Q

What are the metals in group 2 of the periodic table?
EX:
Be- Beryllium (4)
Mg- Magnesium (12) +2
Ca- Calcium (20) +2
Sr- Strontium (38)
Ba- Barium (56)
Ra- Radium (88)

A

Alkaline metals

233
Q

Which group on the periodic table of elements represents the transitional metals?

A

Group 3B to group 12B
Includes popular elements such as:
Ti- Titanium
Cr- Chromium
Mn- Manganese
Fe- Iron
Co- Cobalt
Ni- nickel
Cu- copper
Zn- Zinc
AG- silver
W- tungsten
Pt- platinum
Au- Gold
Hg- mercury

234
Q

What is the property of elements from Group 3A, starting with:
Al: aluminum
Ga: Gallium
In: Indium
Tl: Thalium
Group 4A:
Sn: Tin
Pb: lead

A

Other metals

235
Q

What are the elements:
B- Boron
Si- Silicon
Ge- Germanium
As- arsenic
Sb- antimony
Te- tellurium
Po- polonium

A

Metalloids

236
Q

What are the elemnts:
H- Hydrogen
C- Carbon
N- nitrogen
O- oxygen
p- phosphorus
S- sulfur
Se- selenium

A

Nonmetals

237
Q

What are the elements in group 7A?
F- fluorine
Cl- chlorine
Br-bromine
I- iodine
At- astatine

A

halogens (very reactive and form salts with metals). Do not occur in nature as free elements

238
Q

What are the elements in group 8A
He- helium
Ne- neon
Ar- argon
Kr- krypton
Xe- xenon
Rn- Radon

A

Noble gases

239
Q

What is the first row of elements down at the bottom of the elemental chart (i.e., 57-70)

A

Lanthanoids

240
Q

What is the bottom row of elements at the bottom of the periodic table (i.e., 89-102)

A

Actinoids

241
Q

What is the total number of sulfur atoms represented in 3 Ca(SO4)3?

A

Take the coefficient and multiply by the subscript.

242
Q

What type of bond would fluorine, oxygen, and nitrogen form with a hydrogen atom?

A

-Polar, covalent, and hydrogen bonds.
Polar because if hydrogen’s low electronegativity
Covalent because hydrogen atoms easily bond by sharing electrons
Hydrogen because elements easily accept the hydrogen’s electron to help make up their outer shell.

243
Q

Between which two molecules is a phosphodiesterase bond?

A

PO4 and another molecule

244
Q

What bonds make up sugars?

A

Glycosidic linkages

245
Q

A one molar solution of NaCl contains 58.4 grams of solute in one liter of water. How much NaCl is present in 10 mL of this solution?
A. 58.4 mg
B. 584 mg
C. 0.0584 mg
D. 5.84 mg

A

B. 584 mg

To convert grams per L to grams per mL:
58.4 g / 1L divided by 1000 mL / L
Switch the second fraction upside down and multiply = 58.4 g / 1000 mL.
Divided 58.4 by 1000 to get mL by themselves? = 0.0584 g per mL

Multiply solute of grams per 1 mL (0.0584) by 10 to get how much of the solute is present: 0.584g
Because the answers only present mg as an option. multiply 0.584 by 1000.
= 584

246
Q

How many electrons are in the valence electron shell of K+?

A
  1. For K to be in a neutral state with no charge, it will have one electron in it’s orbit. K has a proton mass of 19, which means that there will be 19 electrons in its orbit when neutral.
    To become K+, potassium would lose an electron thereby having 18 electrons and 19 protons.
247
Q

Which of the following corresponds to the conjugate base of H3PO4?
A. H2PO4 ^-
B. HPO4 ^2-
C. PO4 ^3-
D. H4PO4 ^+

A

A. H2PO4 ^-

For an acid-base pair to exist, two steps must occur.
An acid donates a proton OR a hydrogen
A based accepts a proton OR a hydrogen

If H2PO4 ^- was to accept a proton or a hydrogen, what would its product be? : H3PO4

B. If HPO4 ^2- were to accept a proton or a hydrogen, what would its product be? H2PO4 ^-

C. If PO4^3- were to accept a proton or a hydrogen, what would its product be? HPO4 ^2-

D. If H4PO4^+ were to gain a proton or a hydrogen, what would its product be? H5PO4^2+.
D would represent protonation (adding of a proton) to H3PO4

248
Q

Which of the following is true?
A. materials whose atoms have strongly bonded electrons conduct
B. materials whose atoms have loose electrons insulate electricity
C. materials with free electrons block the conduction of energy
D. materials whose atoms feature tightly bound electrons insulate

A

D.
When electrons are strongly bound to an atom they are not released easy, therefor they do not conduct well.
Examples: air, wood, glass, plastic, cotton and ceramics)

Metals and other substances whose atoms have free electrons that can separate from the atoms and move about are electrical conductors.

249
Q

The concentration of hydrogen ions in a neutral solution is:
A. 7
B. 10^7
C. 10^-7
D. 10^-14

A

C. 10^-7

Formula: H+ = 10^ -pH

250
Q

What will adding an acid to a base always yield?

A

A salt and water.

251
Q

Alcohols have which common structure?
A. benzene ring
B. a carbon atom with a double bond to an oxygen atom and a single bond to a hydroxyl group
C. A nitrogen bond which is also bonded to other carbon atoms
D. A hydroxyl group

A

D. a hydroxyl group (oxygen bound to a hydrogen atom)

252
Q

Compounds with nitrogen bonded to other carbon atoms are called what?

A

amines.
Further classified by how many carbons are attached to the nitrogen.

253
Q

A carbon atom with a double bond to an oxygen atom and a single bond to a hydroxyl group is the function group of?

A

Carboxylic acids

254
Q

Benzene rings are examples of?

A

An aromatic hydrocarbon

255
Q

Structure of a primary amine:

A

R———- NH2
primary amines have nitrogen bound to only one carbon (R functional group)
Secondary amines have two carbons attached
Tertiary has three

256
Q

Which of the following is not true about hydrocarbons?
A. they are volatile
B. they have low boiling points
C. nonpolar/london disperson forces hold the hydrocarbon together
D. Most hydrocarbons are solid

A

D.
Hydrocarbons contain only one hydrogen and one carbon.
They are volatile, reactive, and have low boiling/melting points.
London dispersion forces are present in every molecule.
Hydrocarbons tend to be gases

257
Q

100 g of ethanol C2H6O is dissolved in 100 g of water. The final solution has a volume of 0.2 L. What is the density of the resulting solution?
A. 0.5 g/mL
B. 1 g/mL
C. 46 g/mL
D. 40 g/mL

A

B. 1 g/mL

Density equation
Divide mass of the solution by it’s volume.
m/v

The mass is 200 g (100 g ethanol + 100 g water)
volume is 0.2 L or 200 mL
200 g/ 200 mL - 1 g/mL

258
Q

What is the primary difference between ketones and aldehydes?
A. ketones contain two R-groups and a double-bonded oxygen attached to a central carbon atom, while aldehydes contain only one R-group, a hydrogen atom and a double bonded oxygen atom bonded to the central carbon
B. aldehydes contain two R-groups and a double bonded oxygen attached to a central carbon atom, while ketones contain only one R-group, a hydrogen atom, and a double-bonded oxygen atom bonded to the central carbon
C. ketones are polar substances, and aldehydes are nonpolar substances
D. Aldehydes are polar substances and ketones are nonpolar substances.

A

A.
ketones and aldehydes are very similar structurally.
They each have a double bond between carbon and oxygen.
In ketones there are two carbons bonded to the central carbon, while in aldehydes there is a carbon and a hydrogen bound to the central carbon.

They both have polar and nonpolar properties

259
Q

What type of organic reaction occurs when water is released as two functional groups bonded together?

A

Condensation (opposite of hydrolysis)
Water is released when two groups bond together, where is in hydrolysis, water is added to a compound to break it apart

260
Q

What is the release of a massive amount of energy caused by the bombarding and splitting of a nucleus by a neutron called?

A

Nuclear fission

261
Q

Most organic molecules have all of the following except?
A. high solubility in water
B. a tendency to melt
D. covalent bonds
D. high flammability

A

A. high solubility in water.

Organic molecules contain carbon. Examples: hydrocarbons, lipids, proteins

262
Q

Convert from Fahrenheit to Celsius.

A

F - 32
Divide by 1.8
= C

To convert to Kelvin add 273 to C

263
Q

What is the general formula of a carbohydrate?

A

(CH2O)

264
Q

What is sublimation?

A

The transfer from ice to steam (solid into gas)

265
Q

Which law describes the relationship between volume and temperature with constant pressure?

A

Boyle’s Law

266
Q

What is the law that connect particle count to volume?

A

Avogadro’s Law (Mole)

267
Q

What term defines the maximum strength of a stong acid or basic that a buffer can withstand?

A

Buffer capacity

268
Q

What is a functional group made up of an OH and a caronyl group?

A

Carboxyl group. (COOH).
Makes up amino acids and carboxylic acids.
Also known as hydroxyl (single-bonded OH) and carbonyl (c double bonded to O) joined to a single carbon atom

269
Q

Which is not considered an alloy?
A. steel
B. bronze
C. duralumin
D. chromium

A

D chromium.

270
Q

Which is not a type of chemical reaction that is important to the human body?
A. exchange
B. oxidation
C. synthesis
D. decomposition

A

B. oxidation

271
Q

Which is not a type of energy transfer?
A. convection
B. conduction
C. radiation
D. oxidation

A

D. oxidation.
?

272
Q

What name is given to a compound that cannot be converted by regular chemical processes into simpler chemical substances?

A

Element (NOT atom)

273
Q

Which process builds polymers from monomers?

A

Dehydration synthesis. Requires energy

Hydrolysis is reverse process- reverts polymers back to monomers.

274
Q

What type of reactions release energy?

A

exergonic

275
Q

what type of reactions requires energy

A

endothermic/ endogonic
endothermic absorbs heat

276
Q

What type of reaction releases energy in the form of heat?

A

exothermic

277
Q

The pancrease releases sodium bicarbonate to the duodenome to mix with HCL. What type of chemical reaction is this?

A

Neutralization

278
Q

Antacids react with HCL in the stomach to help alleviate heartburn in what type of reaction?

A

acid-base

279
Q

Direction of radial growth within the elemental periodic table?

A

Radial grows from top to bottom. Right to left.

280
Q

A solution that contains the maximum amount of a dissolved solute is said to be:

A

saturated

281
Q

What accurately models the behavior of an ideal gass?

A

pV=nRT
p = pressure
V = volume
n= number of moles of gas
R= gas constant
T = temperature

282
Q

What is the molar mass of calcium if its atomic mass is 40 amu?

A

40 grams/ mole

283
Q

BF3 is known as

A

boron trifluoride

284
Q

C12
The 12 following the carbon represents 12 —— of carbon

A

atoms

285
Q

Bleaches, disinfectants, and salts often contain elements from which nonmetal group?

A

halogens

286
Q

Rank intermolecular forces in decreasing strength:

A

ionic bonding > hydrogen bonding> ion-dipole interactions> dipole-dipole interactions > London dispersion forces

HESI A2 (8 decks)

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