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Elements and Atoms Elements

Elements consist of only one kind of atom and cannot be decomposed into simpler substances.

Our planet is made up of some 90 elements. (Tiny amounts — sometimes only a few atoms — of additional elements have been made in nuclear physics laboratories, but they play no role in our story). Of these 90, only 25 or so are used to build living things.

The table shows the 11 most prevalent elements in the lithosphere (the earth's crust) and in the human body.

Living matter

 uses only a fraction of the elements available to it

 but, as the table shows, the relative proportions of those it does acquire from its surroundings are quite different from the proportions in the environment.

So,

 the composition of living things is not simply a reflection of the elements available to them

For example, hydrogen, carbon, and nitrogen together represent less than 1% of the atoms found in the earth's crust but some 74% of the atoms in living matter.

 one of the properties of life is to take up certain elements that are scarce in the nonliving world and concentrate them within living cells.

Some sea animals accumulate elements like vanadium and iodine within their cells to concentrations a thousand or more times as great as in the surrounding sea water. It has even been proposed that uranium be "mined" from the sea by extracting it from certain algae that can take up uranium from sea water and concentrate it within their cells.

There is still some uncertainty about the exact number of elements required by living things. Some elements, e.g., aluminum, are found in tiny amounts in living tissue, but whether they are playing an essential role or are simply an accidental acquisition (aluminum probably is) is sometimes difficult to determine.

Atoms

Each element is made up of one kind of atom. We can define an atom as the smallest part of an element that can enter into combination with other elements.

Structure of the atom

Elemental composition of the lithosphere and the human body. Each number represents the percent of the

total number of atoms present. For

example, 47 of every 100 atoms found in a representative sample of the lithosphere are oxygen while there are only 19 atoms of carbon in every 10,000 atoms of

lithosphere.

Composition of the Lithosphere

Composition of the Human Body

Oxygen 47 Hydrogen 63

Silicon 28 Oxygen 25.5

Aluminum 7.9 Carbon 9.5

Iron 4.5 Nitrogen 1.4

Calcium 3.5 Calcium 0.31

Sodium 2.5 Phosphorus 0.22 Potassium 2.5 Chlorine 0.03 Magnesium 2.2 Potassium 0.06

Titanium 0.46 Sulfur 0.05

Hydrogen 0.22 Sodium 0.03

Carbon 0.19 Magnesium 0.01

All others <0.1 All others <0.01

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Each atom consists of

 a small, dense, positively-charged nucleus surrounded by

 much lighter, negatively-charged electrons.

The nucleus of the simplest atom, the hydrogen atom (H), consists of

 a single positively-charged proton. Because of its single proton, the atom of hydrogen is assigned an atomic number of 1.

 a single electron.

The charge of the electron is the same magnitude as that of the proton, so the atom as a whole is electrically neutral. Its proton accounts for almost all the weight of the atom.

The nucleus of the atom of the element helium (He) has

 two protons (hence helium has an atomic number of 2) and

 two neutrons. Neutrons have the same weight as protons but no electrical charge.

The helium atom has two electrons so that, once again, the atom as a whole is neutral.

The structure of each of the other kinds of atoms follows the same plan. From Lithium (At. No. = 3) to uranium (At. No. = 92), the atoms of each element can be listed in order of increasing atomic number. There are no gaps in the list. Each element has a unique atomic number and its atoms have one more proton and one more electron than the atoms of the element that precedes it in the list.

Electrons

Electrons are confined to relatively discrete regions around the nucleus. The two electrons of helium, for example, are confined to a spherical zone surrounding the nucleus called the K shell or K energy level.

Lithium (At. No. = 3) has three electrons, two in the K shell and one located farther from the nucleus in the L shell. Being farther away from the opposite (+) charges of the nucleus, this third electron is held less tightly.

Each of the following elements, in order of increasing atomic number, adds one more electron to the L shell until we reach neon (At. No. = 10) which has eight electrons in the L shell.

Sodium places its eleventh electron in a still higher energy level, the M shell.

From sodium to argon, this shell is gradually filled with electrons until, once again, a maximum of eight is reached.

Note that after the K shell with its maximum of two electrons, the maximum number of electrons in any other outermost shell is eight.

As we shall see, the chemical properties of each element are strongly influenced by thenumber of electrons in its outermost energy level (shell).

The electronic structure of an atom plays the major role in its chemistry.

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The pattern of electrons in an atom — especially those in the outermost shell — determines

 the valence of the atom; that is, the ratios with which it interacts with other atoms, and to a large degree,

 the electronegativity of the atom; that is, the strength with which it attracts other electrons.

Elements with the same number of electrons in their outermost shell show similar chemical properties.

Example 1: Fluorine, chlorine, bromine, and iodine each have 7 electrons in their outermost shell. These so- called halogens are also quite similar in their chemical behavior. When dissolved in water, for example, they all produce germicidal solutions.

Example 2: Those elements with 1, 2, or 3 electrons in their outermost shell are the metals.

Example 3: Those elements with 4, 5, 6, or 7 in their outermost shell are the nonmetals.

Example 4: Helium (with its 2), neon, argon, and krypton (each with 8) have "filled" their outermost shells.

They are the so-called inert or "noble" gases. They have no chemistry at all. Under normal conditions they do not interact with other atoms. So, it is the number and arrangement of the electrons in the atoms of an element that establish the chemical behavior of that element.

This is how it works.

The atoms of an element interact with other atoms in such ways and ratios that they can "fill" their outermost shell with 8 electrons (2 for hydrogen). They may do this by

 acquiring more electrons from another atom

 losing electrons to another atom

 sharing electrons with another atom

The number of electrons that an atom must acquire, or lose, or share to reach a stable configuration of 8 (2 for hydrogen) is called its valence.

Hydrogen, lithium, sodium, and potassium atoms all have a single electron in their outermost shell. Fluorine, chlorine, bromine, and iodine atoms all have 7. Any atom of the first group will interact with a single atom of any of the second group forming, HCl, NaCl, KI, etc. The result of all of these interactions is a pair of atoms each with an outermost shell like that of one of the inert gases: 2 for hydrogen, 8 for the others.

The elements with 2 electrons in their outermost shell interact with chlorine and the other halogens to form, e.g., BeCl

2

, MgCl

2

, CaCl

2

. Again, the result is a pair of atoms each with a stable octet of electrons in its outermost shell.

The elements with 3 electrons in their outermost shell will interact with chlorine in a ratio of 1:3, forming BCl

3

, AlCl

3

.

Carbon atoms, with their 4 electrons in the L shell interact with chlorine to form CCl

4

.

Nitrogen, with its 5 outermost electrons, interacts with hydrogen atoms in a ratio of 1:3, forming ammonia (NH

3

).

Oxygen and sulfur, with their 6 outermost electrons react with hydrogen to form water (H

2

O) and hydrogen

sulfide (H

2

S).

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What determines whether a pair of atoms swap or share electrons?

The answer is their relative electronegativities. If two atoms differ greatly in their affinity for electrons; that is, in their electronegativity, then the strongly electronegative atom will take the electron away from the weakly electronegative one.

Example: Na (weakly electronegative) gives up its single electron to an atom of chlorine (strongly

electronegative) to form NaCl. The sodium atom now has only 10 electrons but still 11 protons so there is a net positive charge of one on the atom. Similarly, chlorine now has one more electron than proton so its now has a net negative charge of 1. Electrically charged atoms are called ions. The mutual attraction of opposite electrical charges holds the ions together by ionic bonds.

Example: Carbon and hydrogen are both only weakly electronegative so neither can remove electrons from the other. Instead they achieve a stable configuration by sharing their outermost electrons forming covalent

bonds of CH

4

. Isotopes

The number of protons in the nucleus of its atoms, which is its atomic number, defines each element.

However, the nuclei of a given element may have varying numbers of neutrons. Because neutrons have weight (about the same as that of protons), such atoms differ in the atomic weight.

Atoms of the same element that differ in their atomic weight are called isotopes.

Atomic weights are expressed in terms of a standard atom: the isotope of carbon that has 6 protons and 6 neutrons in its nucleus. This atom is designated carbon-12or

12

C. It is arbitrarily assigned an atomic weight of 12 daltons (named after John Dalton, the pioneer in the study of atomic weights). Thus a dalton is 1/12 the weight of an atom of

12

C. Both protons and neutrons have weights very close to 1 dalton each. Carbon-12 is the most common isotope of carbon. Carbon-13 (

13

C) with 6 protons and 7 neutrons, and carbon-14 (

14

C) with 6 protons and 8 neutrons are found in much smaller quantities.

Isotopes as "tracers"

One can prepare, for example, a carbon compound used by living things that has many of its normal

12

C atoms replaced by

14

C atoms. Carbon-14 happens to be radioactive. By tracing the fate of radioactivity within the organism, one can learn the normal pathway of this carbon compound in that organism. Thus

14

C serves as an isotopic "label" or "tracer".

The basis of this technique is that the weight of the nucleus of an atom has little or no effect on the chemical properties of that atom. The chemistry of an element and the atoms of which it is made — whatever their atomic weight — is a function of the atomic number of that element. As long as the atom had 6 protons, it is an atom of carbon irrespective of the number of neutrons. Thus while 6 protons and 8 neutrons produce an isotope of carbon,

14

C, 7 protons and 7 neutrons produce a totally-different element, nitrogen-14.

CHEMISTRY-Chapter 2

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MATTER anything that takes up space and has mass

Made of ELEMENTS -can’t be broken down to other substances by chemical reactions • 92 naturally occurring elements

• Each has a unique symbol (usually first one or two letters of name)

• Some symbols dervived from Latin EX: Sodium = Na (from Latin natrium)

• 25 chemical elements are essential to life.

• Four elements—carbon (C), oxygen (O), hydrogen (H), and nitrogen (N) = 96% of living matter • Other 4% of organism’s weight = phosphorus (P), sulfur (S), calcium (Ca), and potassium (K)

• TRACE elements =required in minute quantities

• Some required by all organisms EX: iron (Fe)

• Others only required by some species

EX: humans need 0.15 mg Iodine (I) daily for normal thyroid gland function

ATOMS made of SUBATOMIC PARTICLES

Each kind of atom has a specific number of protons, neutrons, and electrons

• Elements in same row have same # of

electrons in their outer shells

• As move from left to right, one proton & one electron are added to preceding element

• Atoms are electrically NEUTRAL (protons =electrons)

• Atoms that have gained or lost electrons = IONS

= number of protons

= number of protons + neutrons

ATOMIC MASS (1 dalton = 1 amu)

Elements occur in nature as a mixture of ISOTOPES ISOTOPES = atoms with the same number of protons

SUBATOMIC PARTICLE

Electric charge

Mass Location

Proton + 1 dalton In nucleus

Neutron

_

1 dalton In nucleus

Electron 0 negligible Orbit nucleus in energy levels

USES OF RADIOACTIVE ISOTOPES :

Determine age of fossils (carbon dating)

Medical diagnostic and treatment procedures

Research

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but differ in number of neutrons

EX:

14

C ,

13

C, and

12

C all have 6 protons & 6 electrons, but different numbers of neutrons

Most isotopes are stable (EX

: 12

C and

13

C) but some are radioactive (

14

C)

ENERGY LEVELS – 3-D space where electrons are found = ORBITAL

• Level closest to nucleus = lowest energy; outer levels have more energy 1

st

level- 1 orbital holds 2 electrons

2

nd

level- 4 orbitals hold 8 electrons

3

rd

and higher levels- hold increasing numbers of electrons

EX: Lithium (3 ELECTRONS) has two in the first shell; one in second shell Neon (10 electrons) has two in the first shell; eight in second shell

CHEMICAL BEHAVIOR depends on number of electrons in OUTERMOST SHELL (=VALENCE electrons) • Atoms with the same number of valence electrons have similar chemical behaviors

• Atom with a completed valence shell = nonreactive (EX: neon) • Atoms with incomplete valence shells = chemically reactive

• Atoms can give up, accept, or share electrons in order to have a stable outer shell

MOLECULES= two or more atoms of SAME or DIFFERENT elements bonded together (EX: O

2

) COMPOUNDS = two or more DIFFERENT elements bonded together (EX: H

2

O)

CHEMICAL FORMULA = recipe; tells which kinds of atoms and how many EX: H

2

0 = TWO Hydrogen atoms and 1 oxygen atom

*change in characteristics when elements combine = EMERGENT property

TYPES OF BONDS

1) COVALENT: share electrons

• SINGLE- share a PAIR of electrons (shown as single dash H-O-H) • DOUBLE- share TWO PAIRS of electrons (shown as C=C)

• TRIPLE- share THREE PAIRS of electrons (shown as )

POLAR COVALENT BONDS - sharing of electrons = unequal;

seen in atoms with differences in electronegativity

one atom slightly more positive/other more negativity

NONPOLAR COVALENT BONDS- EX: methane electron sharing is equally distributed

2) IONIC BONDS:

electrons are transferred from one atom to another (CATION =+ ANION =-) +/- partners (IONS) are held together by attraction between opposite charges EX: table salt (NaCl) Sodium loses one electron; Chlorine picks up one electron

3) HYDROGEN BONDS: weak attraction between molecules or parts of same molecule • slightly positive hydrogen atom of one molecule attracted to slightly negative atom in another

ELEMENT

# of covalent bonds

Hydrogen 1

Oxygen 2

Carbon 4

Nitrogen 3

Phosphorus 5

Sulfur 2

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EX: water molecule- electrons spend more time orbiting oxygen than hydrogens so oxygen becomes slightly negative and the two hydrogens become slightly positive

4) Van der Waals Interactions-

attractions between ever changing + and - “hot spots” in covalently bonded nonpolar molecules EX: responsible for gecko’s ability to walk up a wall

* Relative strength of bonds: Covalent > Ionic > Hydrogen bond > Van der Waals forces

Individual bonds (ionic, hydrogen, van der Waals) are weak and temporary, but collectively they are strong and play important biological roles.

CHEMICAL REACTIONS- make and break chemical bonds OXIDATION-REDUCTION:

• Oxidation = the loss of electrons (or loss of hydrogen atoms), a molecule that loses an electron is oxidized

• Reduction = the gain of electrons (or gain of hydrogen atoms), a molecule that gains an electron is reduced • Chemical bonds are broken and reformed/atoms are rearranged.

Reactants → products

• Reactions must be “balanced” –Number and kind of atoms in reactants must = those in products

• Matter is conserved in a chemical reaction

• Chemical reactions rearrange matter; they do not create or destroy matter.

• Some chemical reactions go to completion (all the reactants are converted to products)

• Most chemical reactions = reversible (products in forward reaction become reactants in reverse reaction) EX: 3H

2

+ N

2

<=> 2NH

3

hydrogen and nitrogen combine to form ammonia, but ammonia can decompose to hydrogen and nitrogen Initially, reactant concentrations are high, so they frequently collide to create products

As products accumulate, they collide to reform reactants

EQUILIBRIUM

• RATE of formation of products = the RATE of breakdown of products (RATE NOT CONCENTRATION)

• Products and reactants are continually being formed, but no net change in their concentrations

• Concentration of reactants and products typically NOT EQUA; concentrations stabilize at a particular ratio

MOLECULE’S BIOLOGICAL FUNCTION RELATED TO ITS 3-D SHAPE

• Molecule with 2 atoms =linear

• Water molecule is shaped like a V, two covalent bonds are spread apart at 104.5° angle

• Shape of bigger molecule determined by the positions of the electron orbitals shared by bonded atoms

• CARBON- Formation of a covalent bonds leads to hybridization of the orbitals to four new orbitals in a tetrahedral shape

• Large organic molecules contain many carbon atoms with repeating tetrahedral pattern

MOLECULES WIH SIMIALR SHAPES CAN HAVE SIMILAR FUNCTIONS

EX: morphine, heroin, and other opiate drugs = simiilar in shape so they can bind to the same receptors as natural signal molecules called endorphins

Binding of endorphins to receptors on brain cells produces euphoria and relieves pain.

Opiates mimic these natural endorphin effects.

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FUNCTIONAL GROUPS

STRUCTURE

Non-ionized Ionized EXAMPLE CLASS NAME

HYDROXYL

• Polar

• Hydrophilic

• Found in SUGARS

X

ETHANOL GLYCEROL

ALCOHOLS, SUGARS Few AMINO ACIDS

CARBOXYL

• Polar

• weak acid

• hydrophilic

ACETIC ACID

AMINO ACIDS SUGARS FATTY ACIDS

CARBOXYLIC ACIDS.

FATTY ACIDS, AMINO ACIDS

AMINO

• Polar

• Weak base

• hydrophilic

UREA AMINO ACIDS

AMINES AMINO ACIDS

SULFHYDRYL

• Form disulfide bridges

• Help stabilize tertiary structure of proteins

X

Cysteine

THIOLS, DISULFIDE BONDS

PHOSPHATE

• Polar

• Acid

• hydrophilic

• Important in energy transfer

Adenosine triphosphate (ATP) PHOSPHOLIPIDS & DNA

NUCLEOTIDES, PHOSPHOLIPIDS,

ATP

CARBONYL

• Polar

at end of C chain

X

FORMALDEHYDE SUGARS

ALDEHYDE

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• Hydrophilic

in middle of C chain

X

ACETONE SUGARS

KETONE

METHYL

• Non-polar

• Hydrophobic X

FATTY ACIDS OILS WAXES

MEHYLATION OF DNA turns “turns genes

off”

· Each functional group behaves consistently from one organic molecule to another.

· Number and arrangement of functional groups help give molecules their unique properties

EX: TESTOSTERONE (a male sex hormone) and ESTRADIOL (a female sex hormone) are both steroids with same fused four

ring structure but different functional groups attached to the rings result in different functions

References

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