Powerpoint and outline
Chapter 2: Chemistry part 1
Raw materials and fuel for our bodies
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Hi and welcome to the first part of the lecture corresponding to Chapter 2, chemistry.
Learning Objectives
Describe what atoms are, their structure, and how they bond.
Explain water’s features that help it support all life.
Describe carbohydrates—their structure and function.
Describe lipids—their structure and function
Describe proteins—their structure and function.
Describe nucleic acids—their structure and function.
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In this lecture we will cover te following learning objectives <please read the slide>
2.1 Everything is made of atoms.
An element is a substance that cannot be broken down chemically into any other substances.
An atom is a bit of matter that cannot be subdivided any further without losing its essential properties.
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The chemistry that is most important in biology revolves around a few important elements.
A little bit of chemistry goes a long way in the study of biology and in understanding much about your everyday life.
Let’s get started with some basic concepts. What is an element? It is a substance that cannot be broken down chemically into any other substances. And atoms are the smallest bits of matter that cannot be divided any further without losing their essential properties.
Atomic Structure: The nucleus: protons, and neutrons
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All atoms—whether of the element gold or some other element such as oxygen or aluminum or calcium—have the same basic structure.
At the center of an atom is a nucleus, which is usually made up of two types of particles, called protons and neutrons.
Protons are particles that have a positive electrical charge (pro = positive), and neutrons are particles that have no electrical charge.
The amount of matter in a proton or neutron, its mass, is about the same.
Atomic Structure: Electrons
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Whirling in a cloud around the nucleus of every atom are negatively charged particles called electrons.
An electron weighs almost nothing—less than one-twentieth of one percent of the weight of a proton.
Particles that have the same charge repel each other; those with opposite charges are attracted to each other.
Because electrons all have the same charge, the electrons in an atom all repel each other.
But because they are negatively charged, they are attracted to the positively charged protons in the nucleus.
This attraction holds electrons close enough to the nucleus to keep them from flying away, while the energy of their fast movement keeps them from collapsing into the nucleus.
Atomic Number= # of protons
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What distinguishes one element, such as chlorine, from another, such as neon or oxygen?
Atoms of different elements have a different number of protons in the nucleus.
A chlorine atom has 17 protons, a neon atom has 10 protons, and an oxygen atom has 8 protons.
Each element is given a name (and an abbreviation, such as O for oxygen and C for carbon) and an atomic number that corresponds to how many protons it has.
So far, about 90 elements have been discovered that are present in nature and about 25 others can be made in the laboratory.
Everything you see around you is made up of some combination of those naturally occurring elements.
Figure 2-3 The vital statistics of atoms.
Isotopes: same atomic #, different # neutrons
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Atoms don’t always have the exact same number of protons and neutrons in their nucleus and electrons circling around it. They sometimes acquire or lose components.
An atom that has extra neutrons or fewer neutrons than the number of protons is called an isotope. An atom’s charge doesn’t change in an isotope, because neutrons have no electrical charge, but the atom’s mass changes with the loss or addition of another particle in the nucleus.
Carbon, for example, has six protons and so usually has a mass of 12. Occasionally, though, carbon has an extra neutron or two and an atomic mass of 13 or 14. This is rare in nature, but it does occur. These isotopes are called carbon-13 (13C) and carbon-14 (14C) and are referred to as “heavy” carbon. In nature, we frequently see mixtures of several isotopes for a given element. So although a sample of pure carbon is predominantly 12C atoms (with 6 protons and 6 neutrons), there are some 13C and 14C atoms in the sample, too.
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Radioactive Atoms
A few atomic nuclei are not stable and break down spontaneously.
These atoms are radioactive.
They release, at a constant rate, a tiny, high-speed particle carrying a lot of energy.
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Most elements and their isotopes have perfectly stable nuclei that remain unchanged virtually forever, never losing or gaining neutrons, protons, or electrons.
A few atomic nuclei are not so stable, however, and break down spontaneously sometime after they are created. These atoms are radioactive, and in the process of decomposition they release, at a constant rate, a tiny, high-speed particle carrying a lot of energy. (The particle may be a proton, neutron, electron, or just energy.)
For example, uranium-238 (which has 92 protons and 146 neutrons in its nucleus) is a radioactive element. It spontaneously loses a particle containing 2 protons and 2 neutrons, turning it into an isotope of a different element altogether—thorium-236, in this case.
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25 Elements Found in Your Body and the Big 4
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Out of all the elements found on earth, only 25 are found in your body.
The “top 10” most common of these make up 99.9% of your body (Figure 2-5 The body’s chemistry). And just four elements make up more than 96% of your body. These “Big 4” are oxygen, carbon, hydrogen, and nitrogen.
With knowledge about the Big 4, you can understand a huge amount about nutrition and physiology (how your body works), so we’ll focus on the properties of these four elements later in this chapter.
2.2 Electron shells
The chemical characteristics of an atom depend upon number of electrons in their outermost shells.
Atoms are most stable and least likely to bond with other atoms when their outermost electron shell is full.
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While the number of protons identifies an element, it is an atom’s electrons that determine how (and whether) it bonds with other atoms.
Electrons move so quickly that it is impossible to determine, at any given moment, exactly where an electron is.
Electrons are not just moving about haphazardly, though.
Speeding around the nucleus, they tend to stay within a prescribed area.
This area is called an electron shell, and an atom may have several shells, each occupied by its own set of electrons.
Within a shell, the electrons stay far apart because their negative charges repel each other.
Electron Shells
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Electrons move in shells of different energy levels.
First shell has place for 2, second and on for 8
Outermost shell is called a valence shell.
If the outermost shell is full, the atom is stable.
Forming bonds with other atoms may fill the valence shell.
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The first electron shell is closest to the nucleus and can hold two electrons. If an atom has more than two electrons, as most do, the other electrons are arranged in other shells. The second shell is a bit further away from the nucleus and can hold as many as eight electrons.
Atoms become stable when their outermost shell is filled to capacity. When atoms have outer shells with vacancies, they are quite likely to interact with other atoms, giving, taking or sharing electrons to achieve that desirable state of a full outer shell of electrons.
In fact, based on the number of electron vacancies in the outermost shell of an atom, it’s possible to predict how amenable to bonding it will be and even who its likely bonding partners will be.
Why is carbon special?
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Carbon’s electron configuration gives it considerable versatility when it comes to bonding with other atoms and making important compounds
Carbon has 6 electrons overall: 2 in the first electron shell and 4 in the second electron shell. Because the second electron shell has a capacity of 8 electrons, carbon can share its 4 electrons in its outermost shell. This gives it an ability to bond with other atoms in a large number of different ways—including in four different directions—and makes a huge variety of complex molecules possible (see Fig. 2-7).
Carbon most commonly bonds with oxygen, hydrogen, and nitrogen. And because carbon atoms often bind to other carbon atoms, carbon chains are very common—and each carbon bound to two other carbon atoms can still bind with one or two additional atoms. These carbon chains are present in most organic molecules.
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Ions: atoms that gained/lost electrons => charged
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Electrons gained: more negative= anions
Electrons lost: more positive= cations
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Normally, an atom has the same number of electrons as protons.
Sometimes an atom may have one or more extra electrons or may lack one or more electrons relative to the number of protons.
An atom with extra electrons becomes negatively charged, while one lacking one or more electrons is positively charged. Such a charged atom is called an ion
Due to their electrical charge, ions behave very differently from the atoms that give rise to them.
As we’ll see later in the lecture, ions are more likely to interact with other, oppositely charged ions.
2.3 Chemical bonds
Atoms can be bound together in three different ways.
Covalent bonds, in which atoms share electrons, are the strongest.
In ionic bonds one atom loses, the other gains electrons
Hydrogen bonds are weakest and do not form molecules (bond between or within molecules)
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Groups of atoms held together by bonds are called molecules.
It takes a certain amount of energy to break a bond between two atoms. The amount of this energy, called the bond energy, depends on the atoms involved.
In a sense, the molecules are created as short-term storage of energy that can be harnessed later.
There are three principal types of bonds that hold multiple atoms together. <please read the list>
Covalent Bonds
Electrons are shared
One bond: 2 electrons shared
If sharing is equal: nonpolar/apolar bonds (“no poles”)
If sharing is unequal= electrons spend more time closer to one of the atoms: polar molecules (see H2O)
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Covalent bonds are strong bonds formed when two atoms share electrons.
The simplest example of a covalent bond is the bonding of two hydrogen atoms to form H2. Because the atom is most stable when the first shell has two electrons, two hydrogen atoms can each achieve a complete outermost shell by sharing electrons.
The sharing of two electrons among two atoms is called a double bond.
Ions and Ionic Bonds
One atoms loses electrons (ex. Na loses 1 electron, becomes Na ion)
Another takes up electrons (ex. Cl gains 1 electron, becomes Cl ion)
Opposites attract= bond
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Atoms can also bond together without sharing electrons. When one atom transfers one or more of its electrons completely to another, each atom becomes an ion, since each has an unequal number of protons and electrons. The atom gaining electrons becomes negatively charged, while the atom losing electrons becomes positively charged.
An ionic bond occurs when the two oppositely charged ions attract each other and form a compound, a molecule made up of two or more elements.
Ionic bonds are generally similar in strength to covalent bonds. Unlike covalent bonds, in ionic bonds each electron circles around a single nucleus. Ions of equal and opposite charges are attracted to each other and the compound is neutral—that is, it has no charge (Fig. 2-10).
Hydrogen Bonds
Between a H and a negative pole of another molecule (or within a molecule)
Classic example: water molecules
Also present in DNA & proteins
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Ionic and covalent bonds link two or more atoms together.
Hydrogen bonds, on the other hand, are important in bonding multi-atom molecules together. A hydrogen bond is formed between a hydrogen in one molecule and another molecule, often an oxygen or nitrogen atom. This bond is based on the attraction between positive and negative charges.
These atoms are not ions, so where do the electrical charges come from?
The hydrogen atom is already covalently bonded to another atom within the same molecule and shares its electron. That electron will circle both the hydrogen nucleus and the nucleus of the other atom, but the electron is not shared equally. Given that the other atom will always have more than the one proton found in the hydrogen nucleus, the hydrogen electron spends more of its time near the other, more positively charged nucleus than its own nucleus.
Having an extra electron nearby causes the larger atom to be slightly negatively charged, while the hydrogen atom becomes slightly positively charged.
In a sense, the covalently bonded molecules become like a magnet, with distinct positive and negative sides. Magnet-like molecules with distinct positive and negative regions like this are called polar.
2.4 Hydrogen bonds make water cohesive.
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How can certain insects walk on water?
The fishing spider, like numerous other insects such as the water strider, makes use of the fact that water molecules have tremendous cohesion. That is, they stick together with unusual strength. This molecular cohesiveness is due to hydrogen bonds between the water molecules. Each water molecule is “V”-shaped (Figure 2-13 Walking on water!). The hydrogen atoms are at the end of each arm and the oxygen is at the bottom end of the “V” between the two hydrogen atoms.
Oxygen’s strongly, positively charged nucleus pulls the circling electrons toward itself and holds onto them for more than its fair share of the time. Consequently, the oxygen at the bottom of the “V” has a slight negative charge, and the end of the water molecule containing the hydrogen atoms has a slight positive charge.
Hydrogen bonds are much weaker than covalent and ionic bonds and they don’t last very long. Nonetheless, to the fishing spider, the cumulative effect of all the hydrogen bonds in water is to link together all the water molecules in the stream just enough to give the water a surface tension with some net-like properties.
2.5 Water has unusual properties that make it critical to life.
Cohesion=> surface tension, transport in plants
Large heat capacity=> milder climates
Low density as a solid=> ice floats
Good solvent => polar and ionic substances
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All life on earth depends on water; organisms are made up mostly of water and require it more than any other molecule. Hydrogen bonding among water molecules gives water several important properties that contribute to its crucial role in the biology of all organisms:
Cohesion
Large heat capacity
Low density as a solid
Good solvent
Cohesion
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We saw in the previous section how the connection of water molecules through hydrogen bonds makes water cohesive.
The cohesiveness of water molecules also makes it possible for tall trees to exist (Figure 2-14 Like a giant straw).
Leaves need water.
Molecules of water in the leaves are continually lost to the atmosphere through evaporation or are used up in the process of photosynthesis.
In order to get more water, plants must pull it up from the soil.
The problem for many plants, such as giant sequoia trees, is that the soil may be 300 feet below the leaf.
Hydrogen bonds, however, are able to pull up adjacent water molecules to which they have hydrogen bonded.
The chain of linked molecules extends all they way down to the soil, where another water molecule is pulled in via the roots each time a water molecule evaporates from a leaf far above.
Heat Capacity
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It takes a lot of energy to change the temperature of water even a small amount.
The temperature of a substance is a measure of how quickly all of the molecules are moving.
The molecules move more quickly when energy is added in the form of heat.
When we heat water, the added energy doesn’t immediately increase the movement of the individual water molecules. Rather, it disrupts some of the hydrogen bonds between the molecules (As quickly as they can be disrupted, though, they form again somewhere else.
And since the water molecules themselves don’t increase their movement, the temperature doesn’t increase.
The net effect is that even if you release a lot of energy into water, the temperature doesn’t change much.
Because so much of your body is water, you are able to maintain relatively constant body temperatures. Large bodies of water, especially oceans, can absorb huge amounts of heat from the sun during warm times of the year, reducing temperature increases on the land.
Similarly, during cold times of year the ocean slowly cools, giving off heat that reduces the temperature drop on shore.
Low Density as a Solid
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Ice floats.
This is unusual since most substances become more dense when frozen; as the molecules slow down, they pack together more and more efficiently—and densely.
Consequently, they sink.
Water, however, becomes less dense and as you might expect by now, this is due to hydrogen bonding.
As the temperature drops and water molecules slow down, rather than becoming more and more tightly packed, they become less so.
Each V-shaped water molecule bonds with four partners, via hydrogen bonds, forming a crystalline lattice in which the molecules are held slightly farther apart, causing ice to be less dense than water (Figure 2-16 Ice floats).
Good solvent
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If you put a bit of table salt into a glass of water, it will quickly dissolve.
This means that the charged sodium (Na+) and the chloride (Cl) ions that were ionically bonded together all become separated from each other.
The sodium and chloride ions were initially attracted to each other because they are polar molecules, each carrying a slight charge.
Water is able to pry them apart because, as a polar molecule, it too carries charges. Because there is so much salt dissolved in oceans, many of the water molecules have their positively charged sides all facing the Cl ions.
Simultaneously, many molecules of water are turned the other way: with their negatively charged sides facing the Na+ ions.
Consequently, the orderly lattices of hydrogen bonds cannot form in salt water, and it does not freeze well.
The positively charged sodium ions are attracted to the negatively charged side of the water molecule, and the negatively charged chloride ions are attracted to the positively charged side (Figure 2-17 Solutions).
Many substances are polar like water. That is why, like salt, they easily dissolve into it.
Non-polar molecules (such as oil) have neither positively charged regions nor negatively charged regions.
Consequently, the polar water molecules are not attracted to them.
Instead, when oil is poured into a container of water, the water molecules distance themselves from the oil, leaving the oil molecules in isolated aggregations that never dissolve.
2.6 Acids and bases
The pH of a fluid is a measure of how acidic or basic a solution is and depends on the concentration of dissolved H+ ions present.
The lower the pH, the more acidic the solution
Acids, such as vinegar, can donate protons to other chemicals while bases, including baking soda, bind with free protons.
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Let’s talk about acids and bases. What defines an acid such as vinegar or lemon juice? Hydrogen ions dissolved in a solution give it acidity. You may recall that hydrogen is the simplest atom, with only one proton and one electron. That means that a hydrogen ion, that is, a hydrogen atom that lost an electron is basically a proton. The more hydrogen ions in the solution , the more acidic the solution is, and its pH will be lower. Acids donate protons, while bases take them up.
Hydrogen Ions (=proton) and Hydroxide Ions
Ionized Hydroxide Molecule
OH -
Non-Ionized Water Molecule
H2O
O
O
H
H
H
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There’s a lot more going on in water than meets the eye.
Most of the molecules are H2O, but at any instant some of them break up into two parts: H+ and OH.
In pure water, the amount of H+ and OH must be exactly the same, since every time a molecule splits, one of each is produced.
But in some fluids containing other dissolved materials, the fluid can have more H+ or more OH.
pH Scale
The amount of H+ in a solution is a measure of its acidity and is called pH.
pH is the negative log of the H+ concentration, ex. 10-1 (0.1) = pH 1; 10-8 (0.00000001) = pH 8
Acids more H+, lower pH (< 7)
Bases less H+, higher pH (>7)
Neutral = pH7
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The amount of H+ or OH in a fluid gives it some important properties.
In particular, the amount of H+ in a solution is a measure of its acidity and is called pH. The more free hydrogen ions floating around, the more acidic the solution is.
Pure water is in the middle of the pH scale, with a pH of 7.0.
Any fluid with a pH below 7.0 has more H+ ions (and fewer OH ions) and is considered an acid.
Any fluid with a pH above 7.0 has fewer H+ ions (and more OH ions) and is considered a base.
Examples of pH values
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The pH scale is logarithmic, like the Richter Scale for earthquakes: an increase in 1 on the scale represents a ten-fold increase in the hydrogen ion concentration.
An decrease of 2 represents a hundred-fold increase in hydrogen ion concentration.
This means that Coke, with a pH of about 3.0, is 10,000 times (!) more acidic than a glass of water, with a pH of 7.0.
Blood pH
Buffers
can quickly absorb excess H+ ions to keep a solution from becoming too acidic
can quickly release H+ ions to counteract any increases in OH concentration
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The pH of blood is usually 7.4.
Given that most cellular reactions produce or consume H+ molecules, there ought to be great swings in the pH of our blood. Unfortunately, our bodies can’t tolerate such swings. Most of the chemicals that aid in the chemical reactions within our blood or cells stop functioning well if the pH swings by less than half a point.
Fortunately there are some chemicals that act like bank accounts for H+ ions. Called buffers, these chemicals can quickly absorb excess H+ ions to keep a solution from becoming too acidic and they can quickly release H+ ions to counteract any increases in OH concentration.