Unit 1 Flashcards
Which Whmis Symbol is this and what is it for?
Exploding Bomb - For explosion or reactivity hazards
Which Whmis Symbol is this and what is it for?
Flame - For Fire Hazards
Which Whmis Symbol is this and what is it for?
Flame over Circle - For oxidizing hazards
Which Whmis Symbol is this and what is it for?
Gas Cylinder - For Gases Under Pressure
Which Whmis Symbol is this and what is it for?
Corrosion - For corrosive damage to metals, as well as skin & eyes
Which Whmis Symbol is this and what is it for?
Skull & Crossbones - Can cause death or toxicity with short exposures to small amounts
Which Whmis Symbol is this and what is it for?
Health Hazard - May cause or suspected of causing serious health effects
Which Whmis Symbol is this and what is it for?
Exclamation Mark - May cause less serious health effects or damage the ozone layer
Which Whmis Symbol is this and what is it for?
Environment - May cause damage to the aquatic environment
Which Whmis Symbol is this and what is it for?
Biohazardous Infectious Material - For organisms or toxins that can cause diseases in people or animals
What is this and what is it used for? (equipment)
Retort Stand - Used to hold other equipment
What is this and what is it used for? (equipment)
Bunsen burner - Used to heat things
What is this and what is it used for? (equipment)
Wire Gauze - Used to not heat things as intensely
What is this and what is it used for? (equipment)
Test Tube Clamp - Used to hold test tubes for heating purposes
What is this and what is it used for? (equipment)
Evaporating Dish - Leave solutions in them to let them evaporate
What is this and what is it used for? (equipment)
Funnel - To pour liquid into a smaller place
What is this and what is it used for? (equipment)
Scoopula - Scoops or picks up substances
What is this and what is it used for? (equipment)
Test tube holder - Holds test tubes
What is this and what is it used for? (equipment)
Ring Clamp - Used to hold things
What is this and what is it used for? (equipment)
Clay Triangle - Used when heating something really intensely
What is this and what is it used for? (equipment)
Flint Lighter - Lights Bunsen burner
What is this and what is it used for? (equipment)
Burette Clamp - Holds burettes
What is this and what is it used for? (equipment)
Crucible - Gets heated directly
What is this and what is it used for? (equipment)
Spot plate - Used to test for presences of something
What is this and what is it used for? (equipment)
Thermometer & Clamp - Used to hold thermometer and thermometer is used to take temperature of substance. Uncertainty = +/- 0.5 C
What is this and what is it used for? (equipment)
Petri dish - Stores or grow biological substances
What is this and what is it used for? (equipment)
Test tube tongs - Used to pick up test tubes
What is this and what is it used for? (equipment)
Crucible tongs - Used to pick up lid of crucible and crucible
What is this and what is it used for? (equipment)
Erlenmeyer - Used to mix things. Uncertainty = +/- 5%
What is this and what is it used for? (equipment)
Beaker - Measures/holds substances or solutions
What is this and what is it used for? (equipment)
Watch glass - Acts as a lid, to put substances on to measure
What is this and what is it used for? (equipment)
Volumetric Flask - Used to measure things more accurately
What is this and what is it used for? (equipment)
Burette - To add something accurately. Uncertainty = +/- 0.05
What is this and what is it used for? (equipment)
Beaker tongs - To pick up beaker
What is this and what is it used for? (equipment)
Electronic scale - weigh things. Uncertainty = +/- 0.01
What is this and what is it used for? (equipment)
Test tube - Used to heat substances & hold them
What is this and what is it used for? (equipment)
Graduated Cylinder - Used to measure thing precisely. Uncertainty = +/- 0.5 x smallest division
What is this and what is it used for? (equipment)
Pipette pump - used to draw liquid up from pipette
What is this and what is it used for? (equipment)
Pipette - Used to hold a liquid & dispense it
Atom
The smallest part of a substance that cannot be broken down chemically.
Molecule
the smallest particle of a substance that has all of the physical and chemical properties of that substance.
Element
a simple substance that cannot be broken down into smaller parts or changed into another substance.
Solution/Homogeneous Mixture
a mixture in which the composition is uniform throughout the mixture.
Mixture/Impure
made up of two or more pure substances mixed together in any proportion.
Mechanical/Heterogenous Mixture
a mixture in which the composition is not uniform throughout the mixture.
Pure
a single kind of matter that cannot be separated into other kinds of matter by any physical means.
Electron
a negatively charged subatomic particle that can be either bound to an atom or free (not bound).
Proton
A small, positively charged particle of matter found in the atoms of all elements.
Mass Number
the total number of protons and neutrons in an atom.
Neutrons
It exists in the nucleus of the atom alongside protons and makes up an atom’s atomic mass with protons. While a proton is positively charged and an electron is negatively charged, a neutron is neutral; it doesn’t have a charge.
Isotopes
A form of a chemical element in which the atoms have the same number of protons (part of the nucleus of an atom) but with a different number of neutrons (part of the nucleus of an atom).
Average Atomic Mass
the weighted average mass of the atoms in a naturally occurring sample of the element.
Period on Periodic Table
A period is a horizontal row of the periodic table.
Group/Family
The vertical columns on the periodic table
Shell/Orbit
electrons revolve around the nucleus in a specific circular path
Nucleus
The nucleus (plural, nuclei) is a positively charged region at the center of the at
Atomic Structure
The number of protons determines what element it is.
Mass of a proton is 1
Mass of neutron is 1
When adding neutrons you are not changing the atomic but changing the mass number
Mass # - Atomic # = # of Neutrons
# of Neutrons + Atomic # = Mass #
Uncertainties
+/- 0.5 of the smallest division
When adding/subtracting/multiplying/dividing uncertainties make sure you always always add the uncertainties and then do the operation to the other two quantities
Systematic Errors
Systematic errors are due to identified causes and can, in principle, be eliminated. Errors of this type result in measured values that are consistently too high or consistently too low. Systematic errors may be of four kinds:
1. Instrumental. For example, a poorly calibrated instrument such as a thermometer that reads 102 oC when immersed in boiling water and 2 oC when immersed in ice water at atmospheric pressure. Such a thermometer would result in measured values that are consistently too high.
2. Observational. For example, parallax in reading a meter scale.
3. Environmental. For example, an electrical power ‘brown out’ that causes measured currents to be consistently too low.
4. Theoretical. Due to simplification of the model system or approximations in the equations describing it. For example, if your theory says that the temperature of the surrounding will not affect the readings taken when it actually does, then this factor will introduce a source of error.
Random Errors
Random errors are positive and negative fluctuations that cause about one-half of the measurements to be too high and one-half to be too low. Sources of random errors cannot always be identified. Possible sources of random errors are as follows:
1. Observational. For example, errors in judgment of an observer when reading the scale of a measuring device to the smallest division. = UNCERTAINTY
2. Environmental. For example, unpredictable fluctuations in line voltage, temperature, or mechanical vibrations of equipment.
Average atomic mass
Relative atomic mass (means the same thing) = [(% abundance of isotope 1x mass of isotope 1)/100] +
[ (% abundance of isotope 2 x mass of isotope 2)/100] + ……
When talking about trends we always go _____ a Period and _____ a group
Across, down
TREND 1 - Atomic Radius
Atomic Radius: From
Nucleus → Valence Shell
As you go across a period, the atomic radius will decrease because nuclear charge is increasing so the shells are pulled close together & the same shells are constant throughout the period
As you go down a group, the atomic radius will increase because the number of shells increases.
TREND 2 - Ionic Radius
Ionic Radius:
From center of nucleus → Valence shell of Ion
As you go across a period, the ionic radius decreases because there is more pull from the protons until column 14 where there is a jump because of another shell, then decreases again.
As you go down a group, the ionic radius increases because the number of shells increases
TREND 3 - Ionization Energy (IE)
Ionization Energy: Energy required to remove an electron from a valence shell of an atom.
As you go across a period, IE increases because the force of attraction increases because radius gets smaller & number of proton increases
As you go down a group, IE decreases because the force of attraction decrease because radius increases (more shells)
Second Ionization Energy
Xᐩ(g) → X2ᐩ(g) + e-
Energy is needed to remove a second electron from each ion in 1 mole of gaseous 1+ ions. Results in gaseous 2+ ions
TREND 4 - Electron Affinity (EA)
Electron Affinity: Energy required to add an e-
As you go across a period, EA increases because force of attraction increases because number of protons is increase and radius is decreasing
As you go down a group, EA goes down because force of attraction decreases because radius increases
TREND 4A - Electronegativity
Electronegativity: Relative attraction for an e- in a bond
As you go across a period, EA increases because force of attraction increases because number of protons is increase and radius is decreasing
As you go down a group, EA goes down because force of attraction decreases because radius increases
TREND 5 - Reactivity
Metals react by losing electrons to form positive ions
As atomic radius increases first IE will decreases and reactivity will decrease
Non metals react by gaining electrons to form positive ions
As atomic radius increases EN will decrease and reactivity will increase
END
<0.4
95-100%
Covalent Charater
Type of Bonding & Example:
Non-polar covalent
Usually 2 non-metals
eg. H2 = 0
Main/dominant IMF intermolecular force of attraction:
Weak Dispersion Forces
Randerwaals
London Forces
END
0.4-1.7
48 - 95% Covalent
50% Ionic
Type of bonding & Example:
Polar covalent
Between transition metals & non-metals
Or
Metalloids & non-metals
Main/dominant IMF intermolecular force of attraction:
Dipole-Dipole
“H-Bond” (Oxygen, Nitrogen & Fluorine)
Dipole-Ionic
> 1.7
50-100% Ionic
0-50% Covalent
Type of Bonding & Example:
Ionic = metal & non-metal
END = 2.1 (Cl, Na)
Main/dominant IMF intermolecular force of attraction:
Ion-Ion Force of Attraction
Ionic Bonding
Dipole
A bond with oppositely charged ends
Partial Charge
Positive when electronegativity is lower
Negative when electronegativity is higher
Oxidation Number
A way to quantify partial charges
Intermolecular force of attraction
A force of attraction that is between two partices
How many valence electrons does nitrogen have?
5
Consider the electron configuration (arrangement) of boron (2,3) and aluminum (2,8,3). Why are they in the same group (family)?
They are in the same group because they both have 3 valence electrons.
Is a sodium atom larger or smaller than a magnesium atom?
A sodium atom is larger than a magnesium atom. The magnesium atom has an increased nuclear charge and the increased charge attracts the outermost electrons more, causing the atom to become smaller.
Is a sodium atom larger or smaller than a sodium ion?
A sodium atom is larger than a sodium ion because the atom has 3 shells containing electrons and a sodium ion has 2 shells containing electrons.
Based on the trends of the periodic table, what would you predict the charge would be on a stable ion of arsenic?
- 3
Would you expect a neon atom to be larger or smaller than a sodium atom?
A neon atom is smaller than a sodium atom because it has one less shell.
Would you expect a neon atom to be larger or smaller than a sodium ion?
The neon atom would be larger than a sodium ion because they both have the same electron configuration but sodium has an increased nuclear charge (11 protons vs 10 for neon), pulling in the electrons making it smaller.
Explain why beryllium has a higher ionization energy than lithium.
Be has a higher ionization energy because the atom is smaller due to an increased nuclear charge making it more difficult to remove an electron.
What is the most metallic element in the periodic table?
The most metallic element is Fr since it will most readily lose an electron.
What is the most nonmetallic element in the periodic table?
The most non metallic element is F because it will most readily accept an electron.
Which of the following would have the smallest atomic radius?
O2-, F-, Ne, Na+, Mg2+
The magnesium ion would have the smallest atomic radius since it has the greatest nuclear charge (greatest ability to pull outer electrons towards the nucleus resulting in a smaller ion) and all electron configurations are the same of the other elements/ions in the list.
Consider the following electron configurations for neutral atoms:
i) 2,8,2 ii) 2,8,1 iii) 2,8 iv) 2,7 v) 2,5
a) Which of these would you expect to have the lowest ionization energy?
b) Which do you expect to be a noble gas?
c) List the five atoms in order of increasing ionization energy.
a) ii) would have the lowest ionization energy because it has three shells and is a larger atom than i).
b) iii) would be a noble gas since it has a full outer shell.
c) ii, i, v, iv, iii
Using their location of the periodic table, rank each set of elements in terms of increasing atomic size (smallest to largest). Briefly explain the ranking in each case.
Mg, S, Cl
Cl, S, Mg
Mg is the largest atom because it has the lowest nuclear charge in the grouping which creates less of an attraction to the outer electrons, making the atom have a greater radius.
Using their location of the periodic table, rank each set of elements in terms of increasing atomic size (smallest to largest). Briefly explain the ranking in each case.
Al, B, N
N, B, Al
Al is the largest atom because it has 3 shells compared to 2; N is smaller than B because it has more protons in the nucleus (therefore greater nuclear charge density).
Using their location of the periodic table, rank each set of elements in terms of increasing atomic size (smallest to largest). Briefly explain the ranking in each case.
Ne, Ar, Xe
Ne, Ar, Xe
Xe is the largest atom because it has the greatest number of shells.
Using their location of the periodic table, rank each set of elements in terms of increasing atomic size (smallest to largest). Briefly explain the ranking in each case.
Rb. Xe, Te
Xe, Te, Rb
Rb is the largest atom because it has the lowest nuclear charge making the attraction to the outer electrons less, resulting in the largest atomic radius in the grouping.
Using their location in the periodic table, rank each set of elements in terms of decreasing ionization energy. Briefly explain the ranking in each case.
Cl, Br, I
Cl, Br, I - electrons are further away from the nucleus, making them easier to remove
Using their location in the periodic table, rank each set of elements in terms of decreasing ionization energy. Briefly explain the ranking in each case.
Ga, Ge, Se
Se, Ge, Ga - the atomic radius of Ga is the largest, making the outermost electron easier to remove
Using their location in the periodic table, rank each set of elements in terms of decreasing ionization energy. Briefly explain the ranking in each case.
K, Ca, Kr
Kr, Ca, K - the atomic radius of K is the largest, making the outermost electron easier to remove (less attracted to the nucleus)
Using their location in the periodic table, rank each set of elements in terms of decreasing ionization energy. Briefly explain the ranking in each case.
Na, Li, Cs
Li, Na, Cs - Cs has the greatest number of shells, making it lose an outermost electron most easily due to a lesser attraction to the nucleus
Which element in each pair will have the lower electron affinity?
K or Ca
K, since it is a larger atom its nucleus will have a lower attraction to an incoming electron.
Which element in each pair will have the lower electron affinity?
O or Li
Li, since O becomes stable by gaining an electron it has a higher electron affinity.
Which element in each pair will have the lower electron affinity?
S or Se
S has a higher electron affinity since it is a smaller atom than Se and can attract an incoming electron more easily to itself.
Which element in each pair will have the lower electron affinity?
Cs or F
F has a higher electron affinity since it is a non - metal and a very small atom so has the ability to attract an electron to itself.
An element has two isotopes: one of mass 63 u; the other with a mass of 65 u. If the relative abundance of the isotopes is 69.1% and 30.9 %, respectively, find the atomic mass of the element.
63.6
Determine Electronegativity Between two atoms & bond type
Arsenic & Sulfur
0.4 & non-polar covalent
Determine Electronegativity Between two atoms & bond type
Cobalt & Bromine
1.1 & Polar covalent
Determine Electronegativity Between two atoms & bond type
Germanium & Selenium
0.6 & Polar covalent
Determine Electronegativity Between two atoms & bond type
Silicon & Fluorine
2.1 & Ionic
Determine the Oxidation number of the atoms of the specified element
N in NF3
+3
Determine the Oxidation number of the atoms of the specified element
S in S8
0
Determine the Oxidation number of the atoms of the specified element
Cr in CrO4^2-
+6
Determine the Oxidation number of the atoms of the specified element
P in P2O5
+5
Determine the Oxidation number of the atoms of the specified element
C in C12H22O11
0
Determine the Oxidation number of the atoms of the specified element
H in CaH2
-1
Determine the oxidation number of each of the atoms in the compound: H2SO3
H = +1, S = +4, O = -2
Determine the oxidation number of each of the atoms in the compound: OH^-
O = -2, H = +1
Determine the oxidation number of each of the atoms in the compound: HPO4 ^2-
H = +1, P = +5, O = -2
Whats the difference between intermolecular forces and intramolecular forces
Intermolecular forces are between molecules where intramolecular is a bond of molecules
What is the oxidation number of a pure substance e.g O2 or Z
Oxidation number is always zero
Oxygen always has a charge of ___ (oxidation)
-2
Fluorine always has a charge of ___ (oxidation)
-1
When Hydrogen bonds with a non-metal its charge is ___, when hydrogen bonds with a metal its charge is ___
+1, -1
Which group usually has a -1 charge
Halogens
Which group usually has a +1 charge
Alkali Metals
Which group usually has a +2 charge
Alkaline Earth Metals
