Chapter One-Matter, Chemical Trends and Chemical Bonding

Question 0

What is a physical property?

Answer 0

A physical property is a property that you can observe without changing one kind of matter into something new. For example, Iron is a strong metal with a shiny surface, it can be heated to form different shapes, all of these properties can be observed without changing iron into something new.

Question 1

Complete the following…

Matter is anything that has

Answer 1

mass and volume (takes up space)

Question 2

What is a chemical property?

Answer 2

A chemical property is a property that you can observe when one kind of matter is converted into a different type of matter. For example, Iron can react with oxygen to form rust, which is a different kind of matter.

Question 3

What is the difference between a physical and a chemical property?

Answer 3

The difference between a physical and chemical property is that a physical properties can be observed without changing the identity of the substance, while chemical properties can be observed when one kind of matter is converted to another.

Question 4

Fill in the blank. Physical state, colour, odour, crystal shape, malleability, ductility, hardness and brittleness are all examples of ______________________.

Answer 4

Qualitative Physical Properties

Question 5

Fill in the blank. Melting point, boiling point, density, solubility, electrical conductivity and thermal conductivity are all examples of ____________________.

Answer 5

Quantitative Physical Properties

Question 6

Fill in the blank. Reactivity with water, air, dioxide, acids, bases, combustibilty, toxicity and decomposition are all examples of _________________________.

Answer 6

Chemical Properties

Question 7

Properties of matter using words NOT measurements or numerical data.
Ex. Colour, flammability

Answer 7

Qualitative property

Question 8

Properties of matter using measurements or numerical data.

Ex. Density, boiling point

Answer 8

Quantitative property

Question 9

True or False. All non-zero numbers are significant.

Ex. 19.4 = 3 SF

Answer 9

True

Question 10

True or False. All zeros sandwiched between numbers are significant.
Ex. 408 = 3 SF

Answer 10

True

Question 11

True or False. Zeros to the left of the first number are significant.
Ex. 0.0907 = 5 SF

Answer 11

False. Zeros to the left of the first number are NOT significant.
Ex. 0.0907 = 3 SF

Question 12

True or False. Zeros to the right may or may NOT be significant.
Ex. 2000 = 1 SF
2000.0 = 5 SF

Answer 12

True. Decimal adds certainty.

Question 13

True or False. When adding and subtracting, the fewest number of decimal places in the question is the number you should have in your answer.

Answer 13

True.

Question 14

True or False. When multiplying and dividing, the greatest number of SFs in the questions is the number you should have in your answer.

Answer 14

False. When multiplying and dividing, the fewest number of SFs in the questions is the number you should have in your answer.

Question 15

How would you write this number in scientific notation?

0.0000987

Answer 15

9.87x10^-5

Question 16

How would you write this number in scientific notation?

65000000

Answer 16

6.5x10^7

Question 17

Is there another way to perform conversions besides the stair method?

Answer 17

Yes, it is possible to use unitary rates. Using a conversion factor (or unitary rate), you can relate or connect two different units.

Question 18

Define accuracy.

Answer 18

The closeness of a measurement to an accepted value.

Question 19

Define precison.

Answer 19

The closeness of a measurement to other measurements of thee same object or phenomena.

Question 20

Fill in the blanks. Matter can be found in three states; _____________, ____________, ______________. Matter is divided into two categories; _______________________ and __________________. In a __________________, matter is only composed of one type of particle and has _________________________________. An ______________ can be composed of a single atom, like Fe. Or composed of 2 or more of the same atoms, like O2. A _______________ is composed of 2 or more elements joined by chemical bonds, like NaCl. In a ___________________, matter is composed of MORE THAN ONE type of particle. A ______________ is a uniform mixture, like salt water. A ________________________ is a non-uniform mixture (with two or more visible components), like muddy water.

Answer 20

A) Solid G) Unique identifiable properties
B) Liquid H) Element
C) Gas I) Compound
D) Pure Substance J) Mixture
E) Mixture K) Solution
F) Pure Substance L) Heterogeneous Mixture

Question 21

Fill in the blanks. _____________ are the basic substances that make up all matter. The two smallest and least dense of these are _______________ and ____________. Yet, these two account for nearly 98% of the mass of the entire universe!

Answer 21

A) Element
B) Hydrogen
C) Helium

Question 22

In 1809, John Dalton described atoms as solid, indestructible particles that make up all matter. State his atomic theory.

Answer 22

1) All matter is made us of tiny particles (atoms). An atom cannot be created, destroyed or divided into smaller particles.
2) The atoms of one element cannot be converted into atoms of any other element.
3) All atoms of one element have the same properties, such as mass and size. These properties are different from the properties of the atoms of any other element.
4) Atoms of different elements combine in specific proportions to from compounds.

Question 23

True or False.

Particle Symbol Mass (u) Charge
Proton p+ 1 1+
Neutron n 1 0
Electron e- 1 /1837 1-

Answer 23

True.

Question 24

How many _____ are in helium? Protons? Electrons? Neutrons? 4 2 He

Answer 24

Protons = 2 Electrons = 2 Neutrons = 2 Because in scientific notation, the higher number is the mass number, the lower is the atomic number and the letters are the chemical symbol. This is reversed on the periodic table.

Question 25

True or False. Isotopes are atoms of an element that have the same number of protons and the same number of neutrons.

Answer 25

False. Isotopes are atoms of an element that have the same number of protons but different numbers of neutrons.

Question 26

How are ions and isotopes related?

Answer 26

Isotopes and ions have varying numbers of sub-atomic particles. Isotopes have varying neutrons while ions have varying electron values. However, the mass of an ion and an isotope will be different and the mass of an ion and an atom will be the same (the mass of electrons is negligilbe).

Question 27

Fill in the blanks. Some isotopes are more unstable than others. Their nuclei is more likely to decay, releasing energy and subatomic particles in a process called ___________________, which occurs spontaneously. Isotopes that have unstable nuclei are called _____________________.

Answer 27

A) Radioactivity | B) Radioisotopes

Question 28

What is the difference between and isotope and a radioisotope?

Answer 28

An isotope is natural whereas a radioisotope is spontaneous.

Question 29

State the Modern Atomic Theory.

Answer 29

1) Although an atom is divisble, it is still the smallest particle of an element that has the properties and identity of the element. 2) Nuclear reactions (changes that alter the composition of the atomic nucleus) may in fact, convert atoms of one element into atoms of another. 3) Different isotopes of an element have different numbers of neutrons and thus different masses. Scientists treat elements as if their atoms have an average mass. 4) Atoms of different elements combine in specific proportions to form compounds.

Question 30

True or False. Mendeleev sequenced the elements in order of increasing popularity.

Answer 30

False. Mendeleev sequenced the elements in order of increasing atomic mass. Trends began to emerge and this was the creation of the periodic table.

Question 31

There are ________ main groups of elements.

Answer 31

8

Question 32

What is one common trend found throughout the main groups on the periodic table?

Answer 32

Main groups have the same number of valence electrons. Group number = number of valence electrons.

Question 33

Define ionization energy.

Answer 33

The energy that is needed to remove an electron from a neutral atom.

Question 34

Define valence electron.

Answer 34

An electron that occupies the outermost energy level of an atom.

Question 35

Define energy level.

Answer 35

Fixed, 3-dimensional volume in which electrons travel around the nucleus.

Question 36

True or False. Electrons move haphazardly around the nucleus.

Answer 36

False. Their movement is restricted to a fixed region of space called energy shells. An electron moving in a lower energy level is closer to the nucleus.

Question 37

State the Period-Related Pattern. (row)

Answer 37

An element's period number is the same as the number of energy levels that the electrons of its atoms occupy. Ex. period 5 elements have electrons that occupy 5 energy levels.

Question 38

State the Group-Related Pattern. (Column)

Answer 38

You can infer the number of valence electrons in any main-group element from its group number. Ex. Elements in groups 17 (7A) have 7 valence electrons.

Question 39

Fill in the blanks. Lewis structures show ______________ in a _______________ pattern.

Answer 39

A) Valence Electrons | B) Clockwise

Question 40

Why don't noble gases naturally form compounds with other atoms?

Answer 40

There are no electrons available for bonding because of their very stable octet (full outer orbit/energy shell).

Question 41

Describe what happens down each group and across a period based on the periodic trend on atomic size and energy level.

Answer 41

DOWN each group-the size of an atom increases (valence electrons occupy an energy level that is farther and farther away from the nucleus, inner shells block the attraction and repulsive forces between the nucleus and the electrons and the attraction decreases as the radius gets bigger). ACROSS a period-the size of an atom decreases (as you go across a period, proton number and valence electron number increases but energy levels do not, therefore the attractive force via the nucleus pulls valence electrons closer).

Question 42

Fill in the blanks. A charged atom that gains electrons becomes a negatively charged _______________ (non-metals) and an atom that loses electrons becomes a positively charged _____________ (metals).

Answer 42

A) Anion | B) Catoin

Question 43

What are the three main things to keep in mind when examining the relationship between ion formation and the electron arrangement?

Answer 43

1) The energy level from which electrons are gained or given up. 2) The charge on the ion that is formed when an atom gains or gives up electrons. 3) The arrangement of the electrons that remain after electrons are gained or given up.

Question 44

Define isolelectronic.

Answer 44

Same number of electrons. | Ex. Lithium is isoelectronic to Helium.

Question 45

True or False. Ionization energy is the energy it takes to overcome the attractive force of a nucleus to pull an electron away from a neutral atom.

Answer 45

True.

Question 46

Fill in the blanks. The energy required to remove ONE electron from the outer energy level is called the ________________________. It is measured in _________.

Answer 46

A) First ionization energy | B) kJ/mol

Question 47

True or False. Atoms with low ionization energies do not give up electrons easily.

Answer 47

False. Atoms with low ionization energies give up electrons easily.

Question 48

Which group has the lowest ionization energy?

Answer 48

Alkali metals because elements in this group are very reactive because it takes so little energy to remove their single valence electron.

Question 49

Describe what happens DOWN a group and ACROSS each period based on the ionization energy trend.

Answer 49

DOWN each group-ionization energy decreases (easier to remove electrons in higher energy levels, repulsive force and distance from nucleus increases). ACROSS a period-ionization energy increases (attraction between nucleus and valence electrons increases from stronger nuclear pull, number of electrons increases therefore the energy shell is becoming more stable and thus less likely to remove electrons).

Question 50

Define electron affinity.

Answer 50

The measure of the change in energy that occurs when an electron is added to the outer energy level of an atom to form a negative ion.

Question 51

Describe what happens DOWN each group and ACROSS a period based on the electron affinity trend.

Answer 51

DOWN a group-electron affinity decreases (less nuclear attraction thus a smaller change in energy) ACROSS a period-electron affinity increases (number of valence electrons increases resulting in a bigger want for electrons)

Question 52

True or False. When you add an electron, energy is released, therefore electron affinity is high (-).

Answer 52

True. When energy is released, EA is negative.

Question 53

True or False. When you remove an electron, energy is absorbed, electron affinity is low (+) or slightly (-).

Answer 53

True.

Question 54

Fill in the blank. When an element has a high ionization energy it has a ______ electron affinity.

Answer 54

low

Question 55

True or False. Most elements do not exist in nature in their pure form, as elements. Most elements are found in nature as compounds. There are only 90 naturally occurring elements whereas there are thousands of different compounds in natures, and more are constantly being discovered.

Answer 55

True.

Question 56

Define physical properties.

Answer 56

Describing attributes of a substance. | Ex. Appearance, texture, colour, odour, melting point, boiling point, density, solubility and polarity.

Question 57

Fill in the blanks. At room temperature, ionic compounds are found as _______________________. They have a ______ melting point, _________solubility and are ______conductors of electricity when dissolved in water and as a liquid.

Answer 57

A) Crystalline solids B) High C) High D) Good

Question 58

Covalent compounds are found as _________________________ at room temperature. They have a _______ melting point and solubilty. They do not carry electricity well as a ____________ or when ________________________________.

Answer 58

A) Liquids, solids or gases B) Low C) Liquid D) Dissolved in water

Question 59

What is the major difference between an ionic compound and a covalent compound?

Answer 59

Ionic compounds are formed from strong electrostatic interactions between ions (non-metal+metal; transference of electrons). While covalent compounds share electrons between atoms (non-metal+non-metal;sharing of electrons).

Question 60

Fill in the blanks. When 2 atoms ___________ electrons, one atom loses its valence electrons and the other atom gains the electrons. When atoms ____________ electrons, they form an ______________________. Between a _____________________, (__________ low ionization energies and ___________ high electron affinities).

Answer 60

``` A) Exchange B) Exchange C) Ionic bond D) Metal and non-metal E) Metals F) Non-metals ```

Question 61

Fill in the blanks. When atoms __________ electrons, they form a _______________. Between _____________, however it can be between a __________ and a metal with a high ionization energy.

Answer 61

A) Share B) Covalent bond C) Non-metals D)Non-metal

Question 62

How do trends in the periodic table predict chemical bonding?

Answer 62

Electronegativity, a relative measure of an atom's ability to attract shared electrons in a chemical bond. The range of electronegativity can tell us whether the bond is ionic or covalent.

Question 63

True or False. Electronegativity and atomic size trends are inversely related.

Answer 63

True.

Question 64

Why does electronegativity increase will atomic size decreases?

Answer 64

The number of protons in the nucleus increases and the number of full energy shells remains the same, the smaller, the closer and stronger the pull will be. Across a Period, electrons are closer to the nucleus, and thus have a smaller atomic size. The atom attracts a bonding pair of electrons more strongly because the bonding pair can move closer to the nucleus. The atomic size increases and the electronegativity decreases moving down a group because valence electrons are less strongly pulled to the nucleus because the number of filed electron energy levels between the nucleus and the valence shell increases. In a compound, higher energy levels between valence electrons and nucleus means that the nucleus attracts bonding pairs less strongly.

Question 65

True or False. Atomic size, ionization energy, electronegativity and electron affinity are properties of single atoms.

Answer 65

False. Atomic size, ionization energy, and electron affinity are properties of single atoms. In contrast, electronegativity is a property of atoms that are involved in chemical bonding.

Question 66

Why are there no electronegativity values for noble gases?

Answer 66

Noble gases have full valence shells, they are then less reactive and have no desire to form a bond as they are already stable.

Question 67

Fill in the blanks. Electronegativities between 0-0.5 are _____________________ between 0.5-1.7 are ____________________ between 1.7-3.3 are ____________________

Answer 67

A) Pure covalent B) Polar covalent C) Mostly ionic

Question 68

True or False. Atoms with different electronegativities can share electrons unequally without exchanging them.

Answer 68

True.

Question 69

Why are noble gases the most stable elements in the periodic table?

Answer 69

They have full valence orbits (energy levels). When an atom loses, gains or shares electrons through bonding to achieve a filled outer electron energy level, the resulting compound is very stable. Noble gases are extremely unreactive and do not tend to form compounds.

Question 70

What is the octet rule?

Answer 70

Atoms bond in order to achieve an electron configuration that is the same as the electron configuration of a noble gas.

Question 71

True or False. Isoelectronic is when two atoms or ions have the same electron configuration.

Answer 71

True.

Question 72

Why is it called the octet rule?

Answer 72

It is called the octet rule because all the noble gases have 8 electrons in their full outer energy level (except He, which has 2).

Question 73

How does the formation of an ionic bond, for example between sodium and chlorine, reflect the octet rule?

Answer 73

Neutral sodium has 1 valence electron. When it loses this electron to chlorine, the resulting sodium cation has an electron energy level that level contains 8 electrons. Chlorine has an outer electron energy level that contains 7 electrons. When chlorine gains sodium's electron, it becomes an anion that is isoelectronic with argon. Thus, in an ionic bond, electrons are transferred from one atom to another so that they form oppositely charged ions. The strong force of attraction between the oppositely charged ions is what holds them together.

Question 74

What is required for electrical conductivity?

Answer 74

Free to move ions that are negatively and positively charge are needed to carry an electrical current.

Question 75

Fill in the blanks. A liquid has a __________________________ and it is free to move. A solid has an _________________________________________________, so ions cannot move very much. Dissolved states have ______________________ and are free to move.

Answer 75

A) Broken structure, from lattice B) Arranged, rigid lattice formation C) Broken structure, from lattice

Question 76

True or False. Electrons are equally attracted to each atom therefore, instead of transferring electrons; the two atoms die.

Answer 76

False. Electrons are equally attracted to each atom therefore, instead of transferring electrons; the two atoms each share one electron with each other. Each atom contributes to form a covalent bond.

Question 77

Define covalent bond.

Answer 77

Consists of a pair of shared electrons.

Question 78

Why are diatomic elements considered to be pure covalent bond aka non-polar covalent?

Answer 78

Because each element has the same electronegativity therefore they are pure covalent so electrons are being shared equally (no di-pole exists).

Question 79

Do covalent bonds break into ions when the compound melts or boils? Why?

Answer 79

No, whether the compound is in the liquid, solid or gaseous state their bonds do not break, thus they do not break up into ions when they melt or boil. Atoms remain bonded as molecular compounds as pure covalent compounds and do not carry a current.

Question 80

Why can't covalent compounds carry a current?

Answer 80

Covalent compounds cannot carry a current because they do not break into ions (giving them no freedom or room charges which are needed to carry a current). Instead, they remain molecules.

Question 81

What is the difference between intra- and inter- molecular forces?

Answer 81

Intramolecular forces bond atoms to each other within a molecule and intermolecular forces bond molecules to each other. Intramolecular forces are strong covalent bonds and intermolecular forces are weak relative to covalent bonds).

Question 82

Fill in the blank. Covalent compounds have low melting points, due to the ______________________ force being so weak.

Answer 82

Intermolecular.

Question 83

Non-metals tend to form ionic bonds with metals. Non-metals tend to form covalent bonds with other non-metalss. How do metals bond to each other?

Answer 83

In metallic bonding, atoms release their electrons to a shared pool electrons. This forms a non-rigid arrangement of metal ions in a sea of free electrons. They do not have a particular orientation in space.

Question 84

Compare and contrast ionic, covalent, and metallic bonds.

Answer 84

Ionic- electron transer, between a metal and a non-metal Covalent- electron sharing, between a non-metals Metallic bonds-electrons in a shared pool, between metals

Question 85

True or False. With some covalent bonds, the difference in EN is quite significant, but not significant enough for the less electronegative atom to transfer electron(s). The difference is great enough for the bonding electron pair to spend more time near the more electronegative atom than the less electronegative atom.

Answer 85

True.

Question 86

Describe ∆EN 0.0 Pure Covalent Bond

Answer 86

Equal sharing of electrons, no dipoles, nothing than intramolecular forces (limited of very weak intermolecular forces).

Question 87

Describe 0.4 ∆EN Pure Covalent

Answer 87

Slightly unequal sharing of electrons, created dipoles, stronger intermolecular forces (the stronger the dipole, the stronger the intermolecular forces).

Question 88

Describe 1.4 ∆EN Polar Covalent

Answer 88

Unequal sharing of electrons generates negative and positive poles, strong intermolecular forces.

Question 89

Both water and carbon dioxide molecules are considered to be covalent; contain two atoms of the same element bonded to a third atom of another element but both molecules are a different shape. Water is bent while carbon dioxide is linear. Why?

Answer 89

All electrons are involved in bonding carbon dioxide, there are no lone pairs, the linear shape keeps the electrons as far away as possible. To get the same result water molecules are bent.

Question 90

Describe lone pairs.

Answer 90

Electron pairs that are NOT involved in bonding (occur on same atom).

Question 91

Describe bonding pairs.

Answer 91

Electron pairs that are involved in bonding (two different atoms).

Question 92

Electron pairs are arranged around molecules so that they are maximum distance from each other. Why is this?

Answer 92

This is because electrons are negatively charged and they repel each other.

Question 93

Fill in the blanks. Polar molecules have a _______________ and a ______________ charged end. Non-polar molecules do not have ____________ ends.

Answer 93

A) Positive B) Negative C) Charged

Question 94

Fill in the blanks. However, the polarity of a molecule depends on _____________________________________ and _____________________.

Answer 94

A) Presence of polar covalent bonds | B) 3-D shape of molecule

Question 95

What are the rules in determining whether or not a molecule is polar?

Answer 95

1) If ALL the bonds are non-polar than the molecule is non-polar regardless of its shape. 2) If there are polar bonds but the shape is symmetrical such that the polarity of the bonds cancel out, then the molecule is non-polar. 3) If there are polar bonds but no symmetry in the shape, then the molecule is polar.

Question 96

Steps in determining whether or not a molecule is polar.

Answer 96

1) Draw Lewis Diagram 2) Focus in central atom and electron pairs 3) Determine shape 4) Find polarity based on molecular shape