Chemistry Question

A paper analyzing water quality test results based on tables and powerpoint

Background: Below is Water Quality Data from the EBMJG Koi Pond measuring ammonia, nitrate, nitrite, pH, carbonate hardness, and general hardness. These components are regularly tested in different water systems, not just for ponds.

Chapter 5
Water Everywhere: A Most Precious
Resource
•
What are the unique properties of water?
•
Where is the water located that we and other lifeforms use?
•
How does water interact with other chemicals?
•
How do the properties of water change through its interaction with other components?
•
How can we improve the quality of water?
© McGraw Hill
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1
Reflect
Water Everywhere
Watch the chapter opening video and think about the water you drink and use on a daily basis.
a. What substances and impurities are found in this water?
b. Where does this water come from, and where does the wastewater eventually go?
C. How do you think the water habits of a community can affect the natural water supply?
Chapter 5 video
© McGraw Hill
©Tiago Fioreze
2
The Water We Drink
Water is ubiquitous in nature:
• It covers 70% of the Earth’s surface.
• Composes 60% of the human body; 50% of our blood; 77% of the
brain.
Water is essential for life; humans can only go a few days without water.
• Loss of 2% of your body’s water leads to thirst.
• 5% loss gives rise to headaches and fatigue.
• 10 – 15% loss leads to muscle spasms and delusion.
• >15% loss leads to death.
Fresh water is a limited resource!
© McGraw Hill
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The Unique Composition of Water
Water is a liquid at standard temperature and pressure (STP): 25°C and
1 atm
• All other compounds with similar masses are gases under these
conditions (O2, N2, CO2).
Water has an anomalously high boiling point (100°C)
• Liquids with similar molecular structures, such as H2S, have much
lower boiling points.
When water freezes, it expands
• Most other liquids contract when they solidify.
© McGraw Hill
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Electronegativity
The Electronegativity is a measure of the attraction of an atom for
electrons in a chemical bond.
• The greater the electronegativity, the more an atom attracts the
electrons in a bond towards itself.
Table 5.1 Electronegativity Values for Selected Elements
Group 1
2
13
14
15
16
17
H
2.1
18
He
*
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Ne
*
Na
0.9
Mg
1.2
AI
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
Ar
*
*Noble gases rarely (if ever) bond to other elements, and therefore do not have
electronegativity values.
© McGraw Hill
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Polar Covalent Compounds
A difference in the electronegativity of the atoms in a covalent bond
creates a polar covalent bond (a.k.a. polar bond).
• Electrons are not equally shared, but are pulled towards the more
electronegative atom.
• Use arrows to point towards the more electronegative atom; referred
to as a bond dipole.
A nonpolar covalent bond is found between two atoms of the same
element (such as Cl2, O2, N2, etc.).
© McGraw Hill
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Molecular Polarity
A molecule that contains polar bonds may or may not be polar.
• Depends on both the type of bond AND the shape of the molecule.
Water is polar because it has polar bonds and a bent shape.
• The bond dipoles don’t offset or cancel each other.
• BeCl2 is a nonpolar molecule because its polar bonds cancel.
Molecule Polarity
© McGraw Hill
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Hydrogen Bonding
A hydrogen bond is an electrostatic attraction between a hydrogen atom
bonded directly to an atom of N, O, or F and an atom of N, O, or F
1. Hydrogen atom…
2. …bonded to a N, O, or F.
3. N, O, or F in another molecule (could be the same type of
molecule).
Hydrogen bonds are intermolecular bonds
Covalent bonds are intramolecular bonds
© McGraw Hill
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The Properties of Water, Explained
• Hydrogen bonds are not as strong as covalent bonds, but they are
strong enough to affect the physical properties of a substance.
• The high boiling point of water is due to hydrogen bonds, which must
be broken in order to transform water from a liquid to a gas.
• Chemical changes are governed by the strengths of intramolecular
forces (covalent and polar bonds).
• Physical changes are governed by the strengths of intermolecular
forces (hydrogen bonds and London dispersion forces).
© McGraw Hill
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Why Does Ice Float?
• Due to hydrogen bonding,
the structure of ice is porous.
This results in a lower density
for solid water than liquid.
Pipes burst on freezing!
• The solid phase of most
substance is denser than its
liquid.
• Aquatic plants and fish can
live in a freshwater lake
during cold winter because
the lake doesn’t freeze from
the bottom up.
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What Else is Special About Water?
• Specific heat (1.00 cal/g°C) – a lot of energy required to change the
temperature; moist air stores heat energy.
• Heat of fusion – released when the liquid freezes to a solid; spray crops to
prevent freezing.
• Heat of vaporization – released when the gas condenses into a liquid; huge
temperature swing during a thunderstorm.
Energy is required to break the
intermolecular hydrogen bonds
during a phase change
© McGraw Hill
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Fresh Water: A Rare and Precious
Resource!
Only 3% of water found on Earth is freshwater.
•
68% of freshwater is in glaciers, ice caps, snowfields.
•
30% of freshwater is found underground and must be pumped to the surface.
•
0.3% of freshwater is in lakes, rivers, and wetlands.
If all the water on our planet fit into a 2-liter bottle, only 60 mL would be freshwater.
•
Only 4 drops would represent the water in lakes and rivers!
© McGraw Hill
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Water Use Trends
• 322 billion gallons of
water are withdrawn
daily in the US.
• 86% freshwater and
14% saltwater.
• Thermoelectric power
and irrigation
represent the largest
uses of water.
• Agriculture accounts
for 30% of global
water consumption.
© McGraw Hill
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Water Footprints
• A water footprint is an estimate of the volume of freshwater used to
produce particular goods or provide services.
Table 5.2 Water Footprints for Meats
and Grains
Table 5.3 Water Footprints for Various
Products
Food (1 kg)
Water footprint (L,
global average)
corn (maize)
1,200
1 cup of coffee (250 mL)
260
Wheat
1,800
1 cup of tea (250 mL)
27
soybeans
2,100
1 banana (200 g)
160
rice
2,500
1 orange (150 g)
80
chicken
4,300
200
pork
6,000
1 glass of orange juice
(200 mL)
sheep
8,700
1 egg (60 g)
200
beef
15,400
1 chocolate bar (100 g)
1700
b1 cotton T-shirt (250 g)
2500
Product
Water footprint (L,
global average)
Source: Water Footprint Network, 2012
Source: Water Footprint Network, 2012
© McGraw Hill
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Your Turn
2
Differences in Water Footprints
Based on the data in Table 5.2, how do crops compare to
meat, in terms of water usage? What are some reasons for
this?
© McGraw Hill
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Water Pollution
1
• The average American uses
about 100 gallons of water per
day.
• Nearly ¾ of the water entering
our homes goes down the
drain.
• Much of our water comes from
underground aquifers.
© McGraw Hill
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Water Pollution
2
While normally free of pollutants, groundwater can be contaminated by a
number of sources:
• Abandoned mines.
• Runoff from fertilized fields poorly constructed landfills and septic systems.
• Household chemicals poured down the drain or on the ground.
• Manufacturing facilities:
• https://ktla.com/news/local-news/lead-found-in-children-in-communitiesnear-former-exide-battery-plant-in-vernon-usc-study/
• https://www.washingtonpost.com/opinions/2022/09/27/jackson-water-crisisclimate-change-justice/
• https://www.mlive.com/news/flint/2022/09/nrdc-blasts-flint-for-failing-tocomplete-water-service-line-work-by-september-deadline.html
• https://www.theguardian.com/us-news/series/americas-water-crisis
© McGraw Hill
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Solutions
A solution is a homogeneous mixture of uniform composition.
Solutions are made up of solvents and solutes.
• The majority component of a mixture; dissolves the others.
• Minority components of a mixture; dissolved in the solvent.
When water is the solvent, you have an aqueous solution.
© McGraw Hill
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Concentrations of Solutions
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Molarity
Molarity is a commonly used unit of concentration in chemistry
Square brackets [ ] are used to indicate “concentration of” in units of M
[NaCl] = 1.0 M means there is 1.0 moles of NaCl per liter of solution
© McGraw Hill
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Ionic Compounds
• 97% of water on our planet is found in the saltwater of oceans.
• Since water is polar, the partial negative charges on the oxygen
atoms are attracted to the positively charged
ions of the salt crystal.
• Likewise, the partially positive charges on the hydrogen atoms
ions of the salt.
surround
• Dissolving the salt to form its
component ions is called
dissociation.
© McGraw Hill
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© McGraw Hill
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Ionic Compounds with Polyatomic Ions
Ionic compounds with polyatomic ions also dissociate, but the polyatomic
ions remain intact:
Notice the two sodium ions in the compound dissociate from each other
as well, forming a total of 3 separate ions for every unit of Na2SO4 that
dissolves
© McGraw Hill
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Ionic Compounds and Water Solubility
• Not all ionic compounds will dissolve in water.
• Simple generalizations about ionic compounds allow us to predict their
water solubility.
Table 5.5 Water Solubility of Ionic Compounds
Ions
Solubility of
Compounds
Solubility Exceptions
Examples
Group 1 metals, NH4+
all soluble
none
NaNO3 and KBr. Both are soluble.
nitrates
all soluble
none
LiNO3 and Mg(NO3)2. Both are soluble.
chlorides
most soluble
silver, mercury(I), lead(II)
MgCl2 is soluble. AgCl is insoluble.
sulfates
most soluble
strontium, barium, lead(II),
silver(I)
K2SO4 is soluble. BaSO4 is insoluble.
carbonates
mostly insoluble*
Group 1 metals, NH4+
Na2CO3 is soluble. CaCO3 is insoluble.
hydroxides, sulfides
mostly insoluble*
Group 1 metals, NH4+
KOH is soluble. Sr(OH)2 is insoluble.
*Insoluble means that the compounds have extremely low solubilities in water (less than
0.01 M). All compounds have at least a very small solubility in water.
© McGraw Hill
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Ions and Solubility
Ions
Solubility of
Compounds
Solubility Exceptions
Examples
Group 1 metals, NH4+
all soluble
none
NaNO3 and KBr. Both are soluble.
nitrates
all soluble
none
LiNO3 and Mg(NO3)2. Both are soluble.
chlorides
most soluble
silver, mercury(I), lead(II)
MgCl2 is soluble. AgCl is insoluble.
sulfates
most soluble
calcium, strontium, barium,
lead(II), silver(I)
K2SO4 is soluble. BaSO4 is insoluble.
carbonates
mostly insoluble*
Group 1 metals, NH4+
Na2CO3 is soluble. CaCO3 is insoluble.
hydroxides, sulfides
mostly insoluble*
Group 1 metals, NH4+
KOH is soluble. Sr(OH)2 is insoluble.
Name?
© McGraw Hill
Ion
Soluble?
Pb(NO3)2
soluble
CaSO4
insoluble
Na3PO4
soluble
Al(OH)3
insoluble
AgBr
insoluble
Solubility properties are important in
precipitation reactions:
NaCl(aq) + AgNO3(aq)→ AgCl(s) + NaNO3(aq)
Watch the reaction here:

25
Ionic Compounds and Electrolytes
• When ions are in aqueous solutions, the solutions are able to conduct
electricity.
a)
b)
a) Sugar dissolved in water (nonconducting), a nonelectrolyte.
b) NaCl dissolved in water (conducting), an electrolyte.
• The dissociated ions (charge) close the circuit gap in the electrolyte
solution.
• Glowing Pickle Demo: https://www.youtube.com/watch?v=0ijsRfsilQ4
© McGraw Hill
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Strong Versus Weak Electrolytes
• If a compound completely dissociates into ions in water, it is a strong
electrolyte (100% of the substance breaks into its ions).
• If a compound partially dissociates into ions in water, it is a weak electrolyte
(only some of the substance breaks into ions, the rest remains as a whole
uncharged compound).
• If a compound dissolves in water, but does not dissociate into ions, it is a
nonelectrolyte.
• Table sugar, sucrose, a compound that dissolves in water but does not
dissociate.
Rotatable model of sucrose in MolView
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“Like Dissolves Like”
• A polar compound (for example, ethanol) will dissolve in a polar solvent (for
example, water).
• A nonpolar compound (for example, oil) will dissolve in a nonpolar solvent
(for example, gasoline).
• A nonpolar compound will NOT dissolve in a polar solvent, and vice versa.
© McGraw Hill
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Acids
Acids are historically defined as having sour taste, change the color of an
indicator, or react with carbonates.
• Another way to define an acid is as a substance that releases hydrogen ions,
in aqueous solution.
• Since the hydrogen ion has no electron, and only one proton (hence the positive
charge), the hydrogen ion is sometimes referred to as a proton.
Consider hydrochloric acid, dissolved in water:
Since HCl dissociates completely into ions, it is a strong acid.
© McGraw Hill
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The Hydronium Ion
ions are much too reactive to exist alone, so they attach to
something else, such as water molecules.
When dissolved in water, each HCl donates a proton
molecule, forming
• The
to an H2O
the hydronium ion.
remains unchanged:
Hydronium ion – often we simply write
but understand it to mean
when in aqueous solutions.
© McGraw Hill
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Bases
The flip side of the story is the chemical opposite of acids: bases.
A base is any compound that produces hydroxide
ions
in aqueous solution.
Characteristic properties of bases:
• Bitter taste (not recommended).
• Slippery feel when dissolved in water.
• Turns red litmus paper blue.
© McGraw Hill
©McGraw-Hill Education/Eric Misko/Elite Images
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Examples of Bases
Strong bases completely dissociate into
ions in solution
• Examples include Group 1 or Group 2 hydroxides, such as KOH:
• Calcium hydroxide (and other Group 2 hydroxides) produce two equivalents of
What about ammonia (NH3)?
• It is a weak base, even though it has no
• Since this reaction proceeds in both directions, it’s an equilibrium
reaction.
© McGraw Hill
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Neutralization Reactions
When acids and bases react with each other, we call this a
neutralization reaction.
• In neutralization reaction, hydrogen ions from an acid combine with
hydroxide ions from a base to form molecules of water.
• The other product is a salt (ionic compound).
© McGraw Hill
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pH
The pH of a solution is a measure of the concentration of the
ions present in that solution.
The mathematical expression for pH is a log-based scale and is
represented as:
• If
Since pH is a log scale based on 10, a pH change of 1 unit represents a
power of 10 change in
That is, a solution with a pH of 2 has a
ten times that of a solution
with a pH of 3.
© McGraw Hill
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Ion-Product Constant of Water
The three possible aqueous solution situations are:
a neutral solution (pH = 7)
an acidic solution (pH < 7) a basic solution (pH > 7)
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The pH Scale
• The pH scale is useful as it is a measure of acid over many orders of
magnitude
• Tip: The pH is the power of ten of the [H+] without the negative sign – for
example
Acids, Alkalis, and the pH Scale
pH Scale simulation
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Why is Rainwater Naturally Acidic?
• Carbon dioxide in the atmosphere dissolves to a slight extent in water
and reacts with it to produce a slightly acidic solution of carbonic acid.
• The carbonic acid dissociates slightly leading to rain with a pH around
5.3.
© McGraw Hill
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The Chemistry of Acid Rain
• Carbon dioxide is not the only source of
in rain.
• Sulfur oxides (SOx) and nitrogen oxides (NOx) compounds also
dissolve in water forming acids:
sulfuric acid
nitric acid
• This acid rain can wreak havoc
downwind of anthropogenic or natural
sources of SOx and NOx gases.
© McGraw Hill
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Ocean pH
• If rainwater is naturally acidic, why is ocean water basic?
• Three chemical species responsible for maintaining ocean pH:
© McGraw Hill
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Ocean Acidification
1
• Ocean pH is decreasing due to increased atmospheric carbon dioxide.
• Carbonate ions
are necessary for marine animal shells and
skeletons.
•
produced from the dissociation of carbonic acid reacts with
carbonate ion in seawater:
• Calcium carbonate in the shells of sea creatures begins to dissolve
to maintain the concentration of carbonate ions in seawater:
Ocean Acidification: “The Other Carbon Dioxide Problem”
Ocean Acidification and Chemical Signalling
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Ocean Acidification
2
Over the past 200 years, the amount of carbon dioxide in the atmosphere
has increased, so more carbon dioxide is dissolving in the ocean and
forming carbonic acid.
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Aquatic Life and pH
Acidification of waters occurring in lakes and streams, too.
• Midwestern states have considerable limestone (CaCO3) that
neutralizes acid (called acid neutralizing capacity, ANC).
• New England states have largely granite, which is much less reactive
so the lakes and streams are more sensitive.
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