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Chemistry > Matter, Chemical Bonding, Etc > Flashcards

Matter, Chemical Bonding, Etc Flashcards

1
Q

What is chemistry?

A

The study of matter and its changes and properties

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2
Q

What is matter?

A

Anything that has mass and takes up space

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3
Q

What are the two types of matter and the two groups within each?

A

Mixtures: heterogeneous/mechanical mixtures and homogeneous/solutions
Pure substances: elements and compounds

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4
Q

Explain heterogeneous mixtures

A

2 or more components with two or more visible phases. Variable composition. Can be separated by physical means. (Ex. Pizza)
Also: suspensions (ketchup, paint) and Colloids (particles can’t be seen with naked eye, milk)

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5
Q

Explain homogeneous mixtures.

A

Two or more components. Only see one phase. Variable composition. Can be separated by physical means. Ex. Salt water, air, alloids

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6
Q

Explain elements

A

Simplest pure substance. Can’t be decomposed. Found in periodic table. Represented with symbols.

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7
Q

Explain compounds

A

Two or more elements. Constant composition. Decomposed by chemical means. Represented by formulas. Smallest particles are molecules.

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8
Q

Give examples of physical properties.

A

The appearance: texture, colour, state, clarity, hardness, melting point, boiling point

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9
Q

Five proofs of chemical change

A
  1. Gas is released (fizzing, bubbling, foaming)
  2. Energy is released or absorbed
  3. Precipitate formation
  4. True colour change
  5. Hard to reverse
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10
Q

What did Democritus propose?

A

First atomic theory. If matter is divided into smaller and smaller pieces, there will eventually be a part that can no longer be divided (the atom)

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11
Q

What did john dalton contribute to the atomic theory?

A

Billiard ball model. All ,after consists of particles called atoms. Atoms of each element are unique. Atoms are undivisible.

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12
Q

What did JJ Thompson contribute?

A

Cathode ray tubes. Raisin bun model.

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13
Q

Rutherford’ contribution

A

Gold foil. Conclusion: matter consists of mostly empty space with a positive nucleus

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14
Q

Explain how to find protons, neutrons, electrons

A

Atomic number = e and p

N = mass number -protons

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15
Q

List the maximum number of electrons in each level (s, p, d, f)

A

1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14

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16
Q

What letter describes each thing: energy level, shape, orientation, spin)

A

Energy level is n
Shape is l
Orientation is m
Spin is S

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17
Q

What is an isotope?

A

Atoms of an element that contain the same number of protons but different number of neutrons. Therefore have different masses.

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18
Q

Describe unstable isotopes

A

Have excess nuclear energy. Radioactive. Emit radiation. Called radioisotopes.

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19
Q

Formula for average atomic mass

A

AAM = (abundance as decimal)1(mass)1 + (abundance as decimal)2(mass)2

20
Q

Explain the different periodic trends.

A

Atomic size/radius: size of an atom from its nucleus to where electrons spend 90% of their time)
increases as you move down the groups because there is an increase in energy levels, there is less attraction to the nucleus with more energy levels so more spread out
Decreases across a period because nuclear charge increases and so does attraction towards the nucleus.

Ionization energy: the amount of energy required to remove an electron.
Decrease down groups, due to the larger energy levels and less attraction, less energy is required to remove them
Increase across periods since the attraction to the centre increases as the number of protons in nucleus does.

Electron affinity: the liking for electrons (how willing it is to take or share them)
Decreases down the groups as the higher energy levels makes the attraction of electrons lower.
Increases across periods since the positive charge in the centre is increasing.

Electronegativity: willingness to share electrons.
Same as above.

21
Q

Explain bonding.

A

Atoms bond to achieve an electron configuration identical to that of the nearest noble gas (isoelectronic). This is to reach a full, complete or stable octet.

22
Q

Explain the different types of bonding

A

Covalent: takes place between two or more non metals. Electrons are shared
Ionic: takes place between metals and non metals. Electrons are transferred.

23
Q

List diatomic molecules and what they are.

A

Two identical atoms bonded together.

N2, O2, N2, F2, Cl2, Br2, I2

24
Q

Explain the different shapes of molecules

A

Linear, trigonal-planar (three), tetrahedral (four), trigonal-bypyramidal (5), octahedral (6)

25
Q

Explain the development of the atomic theory.

A
  1. Greek philosophers
  2. Democritus (breaking down, etc)
  3. Alchemists (not interested in matter, just wanted to turn lead to gold)
  4. J Dalton (billiard ball)
  5. JJ Thompson (raison bun model, cathode ray tubes)
  6. Ernest Rutherford (gold foil and planetary model)
  7. Niels Bohr (electrons only exist in specific energy levels)
  8. Erwin Schrodinger henergy clouds)
26
Q

Explain the development of the periodic table

A
  1. Elements were ordered by mass
  2. J. Dorbereiner noticed some sets of three elements were similar. Organized them into triads.
  3. J. Newlands arranged elements in order of increasing atomic mass. Every eight elements formed a new group. (Problem: he thought all elements had been discovered.
  4. Demetri Mendeleev Arranged them in order of their increasing atomic mass and periodicity. Left space for unknown elements and predicted the characteristics of those elements. Didn’t predict the finding of the noble gases.
27
Q

Explain the classical system of naming ionic compounds when the metal has two possible valences.

A 
ic (high) and ous (low) can be used as endings to indicate the valence. 
Iron: ferr 
Copper: cupr 
Gold: aur 
Lead: plumb
Tin: stann 
Other metals: use regular names.
28
Q

Explain how to name compounds when they contain hydrogen.

A

If H is the cation, hydrogen is used first and then anion’s end turns to ide.
If H is the anion, it changes to hydride.

29
Q

List their common names:
NH3
PH3
SbH3

A

Ammonia, phosphine, stibine

30
Q

Explain the different rules for naming polyatomic ions.

A

Base is ate.
One less oxygen becomes ite.
2 less becomes hypo __ ite
One more becomes per __ ate

31
Q

State the formulas for each:

Chlorare, bromate, iodate, nitrate, carbonate, chromate, sulfate.

A

ClO3, BrO3, IO3, NO3, CO3, CrO4, SO4

32
Q

Explain how to name binary acids.

A

These are acids that contain hydrogen and a metal. It indicate that these are acids, they will have aq at the end. You write hydro__ic acid.
Ex. HF is hydrogen fluoride.
HF (aq) is hydrofluoric acid.

33
Q

What are oxyacids?

A

Acids that contain an oxygen iron in addition to a hydrogen ion an another element. During reactions with these substances, water and a salt are produced.

34
Q

Explain how to name oxyacids.

A

If they are ate ions, then it will be ic.
Example: h2SO4 is sulphuric acid.
If they are ite ions, then it will be ous.
Ex. H2SO3 is sulphurous acid.

35
Q

What is an acid salt?

A

An acid salt is the compound formed when only some of the hydrogen atoms in
an acid, are replaced by a metal.

36
Q

Explain how to name acid salts.

A

List metal, hydrogen, then polyatomic ion. If there are two hydrogen’s, say di in front.

37
Q

What is a hydrate?

A

When a salt has water molecules attached to it.

38
Q

Explain how to name hydrates.

A

Write compound before period. Then write water with prefix in front of it.

39
Q

Explain metallic bonding.

A

Electrons are shared in a pool.
Malleability: flatten into sheets by hammering electrons to the side
Ductility: stretch into wires.

40
Q

Explain what a melting point and boiling point is.

A

The temperature at which a solid turns to a liquid at standard atmospheric pressure.
The temperature at which a liquid turns to a gas at basic atmospheric pressure.

41
Q

Explain how melting works.

A

There is always kinetic energy in a solid, even though they are tightly packed together. When energy is sent into a solid through heat, the kinetic energy keeps increasing, as well as the heat. Once the kinetic energy gets high enough for the particles to break apart, the temperature will stop increasing and the substance will melt.

42
Q

Explain how boiling works.

A

In,iquids, particles have enough kinetic energy to slide past one another. Once the substance reaches its boiling point, the particles will pull completely away from each other and bounce off each other as a gas.

43
Q

Explain what it means for a substance to have a high, low or intermediate boiling point.

A

High: lots of energy is required and the forces must be very strong between them. Substance is therefore an ionic compound.
Intermediate: polar molecular compound. The partially positive/negative ends have somewhat of an attraction, but don’t require as much energy to pull apart.
Low: non-polar molecular compound. Not a lot of energy Is required.

44
Q

Explain solubility.

A

For something to dissolve in water, the slightly positive end of the water molecule must be more attracted to the negative ion or slightly negative end of another compound than to itself. Non-polar molecules are not soluble.

45
Q

Explain electrical conductivity.

A

Electrons are needed to move freely in order to conduct an electric current. Ionic compounds can only conduct electricity when dissolved in water. Molecular compound cannot conduct electricity regardless of their polarity, as the electrons are never freely moving. Only acids, as they become ionic when dissolved in water.

46
Q

Eplain dipole to dipole attraction.

A

When the partially positive end of one dipole is attracted to the partially negative end of another.

47
Q

Explain intermolecular and intramolecular forces.

A

An intramolecular force is any force that binds together the atoms making up a molecule or compound, not to be confused with intermolecular forces, which are the forces present between molecules.

Chemistry (3 decks)

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