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The smallest particles of matter are atoms. Atoms have a nucleus, with protons and neutrons as major components and electrons which orbit the nucleus.

• A compound, composed of the same kind of atoms is called an element. Each element has a characteristic number of positively charged protons in its nucleus.
• The number of protons present in a nucleus is called atomic number.
• The total number of protons and neutrons in an nucleus are called the mass number.

The Mass of Atoms

The mass of atoms is very small in the order of 10^-23 to 10^-22 g. Because the use of such small numbers is inconvenient, a relative mass scale is used, known as atomic mass units (abbreviated amu) or sometimes called dalton. The scale is based on the carbon-12 atom which has 6 protons, 6 neutrons and 6 electrons and and has a mass of 1.9926 x 10^-23g. Thus:

• 1 amu = 1/12 of the mass of a C-12 atom = 1.66054 x 10^-24 g
• The mass of a proton is 1.6726 x 10^-24g = 1.00728 amu
• The mass of a neutron is 1.6749 x 10^-24g = 1.00867 amu
• The mass of a an electron 1/1837 of the proton mass = 0.000549 amu

Using these numbers we can calculate the relative atomic masses of a hydrogen atom and helium atom:

• atomic mass _H = 1.00728 amu (mass of proton) + 0.00055 amu (mass of electron) = 1.00783 amu.
• atomic mass_He = 2 x 1.00728 + 2 x 1.00867 + 2 x 0.00055 = 4.03300 amu

However, these values are not the actual atomic masses of the H and He atoms. When a free proton and a free neutron combine with each other to form a hydrogen atom (or 2 protons, 2 neutrons and 2 electrons in case of He), some small fraction of the mass is lost in form of released energy . This is called the mass defect.

calculated atomic mass of He atom (amu)	actual atomic mass of He atom (amu)	mass difference (amu)
4.03300	4.00260	0.03040

The amount of energy released is equivalent to the mass difference of 0.0304 amu according to Einstein's formula E = m c² . The larger the nuclei, the more mass is lost in form of energy.

Isotopes

If you check atomic masses in the periodic system you will find quite often that the value is higher than you would have expected from the sum of proton -, neutron-, and electron - masses minus mass defect. For example chlorine has 17 protons and assuming that there are 18 neutrons, the atomic mass of chlorine should be approximately 35 amu (you could calculate the exact expected value, neglecting the mass defect or mass loss caused by the formation of the nucleus) However, the actual atomic mass of chlorine is 35.4527 amu. The reason for this is that chlorine occurs in nature in two different forms. One form has 17 protons and 18 neutrons, the other form has 17 protons and 20 neutrons. Atoms with the same number of protons but different number of neutrons are called isotopes.

Naturally occuring carbon consists to 99% of the carbon-12 isotope (6 protons and 6 neutrons) 1% of the carbon-13 isotope (6 protons and 7 neutrons) and a trace of the unstable carbon-14 isotope (6 protons and 8 neutrons). Concequently the atomic mass of naturally occuring carbon is not 12.00000 amu as it would be for the pure C-12 isotope but 12.01115 amu.

Atomic Mass or Atomic Weight

The atomic mass or atomic weight - as it is more frequently called - of an atom is the average mass of the mixture of isotopes that reflects the masses and relative abundance of the elements as they occur in nature. For example, for hydrogen we would expect :

• mass_H-atom = 1.00728 amu (proton) + 0.00055 amu (electron) - 0.00004 amu (mass defect) = 1.00779 amu

	natural abundance	atomic mass (amu)	Calculation
H atom	99.985 %	1.00779	1.00728 amu (proton) + 0.00055 amu (electron) - 0.00004 amu (mass defect) = 1.00779
D atom	0.015 %	2.01355	1.00728 amu (proton) + 1.00867 amu (neutron) + 0.00055 amu (electron) - 0.00295 amu (mass defect) = 2.01355
naturally occurring Hydrogen		1.00794	1.00779 amu x 0.99985 + 2.01355 amu x 0.00015 = 1.00794 amu

additional examples for the calculation of the atomic mass of Mg.

• The atomic masses or atomic weights you see listed for each element in the periodic system are weighted averages of the masses (measured in amu) of all natural isotopes of the element

Radioactivity

Most elements exist in several isotopic forms. Some isotopes are instable and their nuclei break apart. During this break up, energy is emitted in form of radiation and the element is said to be radioactive.

Radiation

There are 3 different types of radiations

• Alpha Radiation
  The radioactive nucleus becomes more stable by giving off alpha radiation. Alpha radiation consists of helium nuclei ( no electrons). For example
  U-238 emits alpha radiation to form Th-234.
• Beta Radiation
  Beta radiation consists of electrons. During beta decay a neutron in the nucleus is converted into a proton and an electron. The electron is emitted. The carbon-14 isotope is a beta emitter.
• Gamma Radiation
  Gamma radiation is a high energy electromagnetic radiation similar to X-rays. Gamma radiation is usually associated with alpha or beta decay. For example
  Ra-226 decays to Rn 222 + He (alpha radiation) + gamma radiation

Half Life

The rate of radioactive emission or the rate of nuclear decay is measured in half life. The half life time of an radioactive element is the time it takes for half of a given amount of an radioactive element to degrade. Using the radioactive decay equations one can calculate the decreasing amounts of a radioactive element over time. For example Radium-226 has a half life time of 1600 years. That means 1g of Ra-226 will have degraded to 0.5 g after 1600 years.

• Carbon-14 is often used to date organic matter. With a half life 5730 years C-14 dating is used to date objects as old as 40000 years.
• Uranium-235 has a half life time of 7.04 x 10⁶ years. It decays to a number of intermediary isotopes, with Pb-207 being the final stable isotope . Using ratios of the intermediate isotopes and Pb-207, rocks can be dated to the beginning of the formation of the earth (4.8 billion years ago).
• P-32 a beta emitter is mostly used in molecular biology research. Its half life is 14 days.
• I-131 (half life = 8 days, beta and gamma) and I-123 (half life 13 hrs) can be used to test the thyroid function.

The first model of atoms by Bohr, depictures electrons orbiting the nucleus at discrete distances like planets circling the sun. Quantum mechanics has refined this model. Heisenberg has pointed out that it is impossible to determine accurately the position and speed of the electron. Thus, the electrons position around the atom has to be described as a space in which there is a high probability to find the electron. This region is called an orbital. Atoms have many orbitals, they differ in energy, shape and their location with respect to the nucleus.

Electrons in orbitals are grouped into certain principal energy levels or shells and are designated n = 1, 2, 3, 4, 5, ... The maximum number of electrons in a shell is 2n². Each shell has subshell energy levels which can be filled with electrons. The subshells are designed : s, p, d, f, g, h, i, with a s-orbital having the lowest energy. The subshell orbitals have a distinct shape. A s-orbital has a spherical shape and holds a maximum of 2 electrons with opposite spin

Maximum number of electrons per shell

n=1	1s2
n=2	2s2	2p6
n=3	3s2	3p6	3d10
n=4	4s2	4p6	4d10	4f14
n=5	5s2	5p6	5d10	5f14	5g18

The subshells are filled with electron in the order of increasing orbital energy.
The orbitals are filled with electrons starting with the lowest energy orbital (1s), up to the highest energy orbital shown here (7p). The enegy of orbitals within a row is increasing, however, the energy difference between orbitals within the row is relatively small. The energy difference between orbitals in different rows is large, e.g. the energy difference between 4p and 3d is small, however, there is a big increase in energy from 4p to 5s.
Xe has the following electron configuration:
1s², 2s², 2p⁶, 3s², 3p⁶, 4s², 3d¹⁰, 4p⁶, 5s², 4d¹⁰, 5p⁶

The chemical properties of elements show a pattern which repeats itself in in periods. This periodicy, which results from the way electrons are filled into the s, p, d and f orbitals is the basis for arranging the elements in the periodic table. Horizontal rows are called periods, columns are called groups.

Main group Elements of the main group IA (or group 1) called alkali metals and IIA (or group 2) elements called alkaline earth metals fill up their s-orbitals. The remaining elements of the main groups IIIA (or group 13) - VIIIA (or group 18) fill up their p-orbitals. Transition elements or B-group elements fill up their inner shells having no more than two electrons in their valence shell. Transition metals of groups IIIB (or group 3) to VIIIB (or group 10) and groups IB (or group 11) and IIB (or group 12) fill up their d-orbitals and the lanthanides and actinites add electrons to their f-orbitals. Members of the VIIIA group (or group 18), the noble gases have their outer shells completely filled with 8 electrons. This stable arrangement is the reason for the chemical inertness of these elements. The tendency to attain 8 electrons in the outer shell is called the octet rule. Elements can achieve a stable octet electron configuration by gaining or losing electrons to form ions or by sharing electrons to form covalent bonds.

Ions

Cations (positively charged ions) are formed if electrons are removed from an atom. The energy necessary to remove one electron from an atom is called ionization energy. For example the energy required to remove one electron from a sodium atom is:
Na ® Na⁺ + e^-          E_ionization = 2 x 10^-19 cal or 5.1 eV .
Electrons whose orbitals are close to the nucleus are more difficult to remove and have a higher ionization energy. Electrons in higher orbitals have more energy and are easier to be removed, thus having lower ionization energies.
For example, Be: removal of a 2s electron = 9.3 eV compared to Ba: removal of a 6s electron = 5.2 eV.

The tendency of an atom to attract electrons is called electronegativity. If the electronegativity is strong enough, the electron can be transferred completely to the atom, forming a negatively charged ion or anion. The electronegativity increases within a period ( e.g. from carbon to fluorine) but decrease within a group. For example O is more electronegative than N but less electronegative than F, but O is more electronegative than S, Se or Te.

Atomic and Molecular Interactions

The outer electrons of atoms can interact to form covalent bonds , ionic bonds. In addition molecules can form non covalent associations such as polar and non polar interactions and so called hydrogen bonds.

Covalent bonds.

IA	IVA	VA	VIA	VIIA
H
	C	N	O	F	All non metallic elements listed to the left can form covalent bonds. The semiconductor elements B, Si, Ge, As, Sb, Te, Po and At can also form covalent bonds.
		P	S	Cl
			Se	Br
				I

Single covalent bonds are formed if each atom contributes one electron to the bond. The driving force behind bond formation is to reach noble gas conformation . For example hydrogen has one electron, two hydrogen atoms can form a covalent bond and by sharing the two electrons they can attain the electron configuration of He (2s²)
where the horizontal line represents 2 electrons.
Hydrogen can also form a covalent bond with fluorine. F has the 7 electrons in its outer shell (2s²,2p⁵). By sharing the electron from one hydrogen, fluorine can reach the octet configuration

Oxygen has an outer shell configuration of 2s², 2p⁴. To attain a stable octet electron configuration it needs to share 2 electrons from hydrogen. Since H has only one electron, oxygen will bind two hydrogens to form H₂O. Because of the same reasons nitrogen needs three electrons to form NH₃.

Judging from its outer shell electron configuration 2s², 2p² carbon should be two valent, since it has 2 single electrons occupying the two p-orbitals. However, carbon forms compounds using four energetically equivalent orbitals. This is achieved by combining one s-orbital and three p-orbital to obtain four sp³ hybrid orbital. This process is called hybridization . In case of carbon the four sp³ hybrid orbitals are arranged tetrahedrally, a shape in which the four negatively charged electrons are most distant from each other. Carbon will bond four hydrogen atoms and by sharing the electrons it can attain octet conformation. Other atoms such as oxygen and nitrogen can also undergo hybridization.

A double bond is formed when each atom contributes two electrons to a bond. When carbon forms double bonds it uses the 2s² and two 2p¹ orbitals to form three sp² hybrid orbitals , with one electron remaining in the 2p_y orbital. The resulting bond angles of the three sp² hybrid orbitals (in grey color) are arranged at 120^o as shown below for the compound ethene (H₂C=CH₂).

Triple bonds are formed when a total of 6 electrons are shared between two atoms. Nitrogen has 2 paired electrons ( | ) in its 2s orbital and three electrons in its 2p orbitals. These 3 electrons can bond with with 3 electrons form another nirogen atom to form N₂

When carbon forms triple bonds it undergoes a sp hybridization i.e. only one p-orbital is combined with the s-orbital, while remaining two electrons are placed in a 2p_y and a 2p_z orbital. The bonding orbitals for the compound acetylene is shown below.

Carbon is also sp hybridized in CO₂ forming two double bonds.

Covalent bonds are strongest atomic interactions. Bond energies are between 60 - 110 kcal/mol for single bonds, approximately 150 kcal/mol for double bonds and about 200 kcal/mol for triple bonds

Ionic bonds
If an electron is transferred form one atom to another atom two opositely charged atoms or ions are formed. The force of attraction between the ions is called an ionic bond. All metals are capable of forming ionic bonds. For example Na has one electron (3s¹) in its outer shell and Cl has 7 electrons (3s², 3p ⁵). Na can attain octet conformation if it loses one electron, whereas Cl has to gain one electron to reach noble gas configuration forming : Na⁺ Cl^-

Bond energies for ionic bonds range from 4 - 8 kcal/mol and are thus much weaker than bond energies of covalent bonds.

Polar and non polar interactions
The distribution of the shared electrons in covalent bonds can be symmetrical if the electronegativity of both atom is similar as for example in O=O, H-H or C-H bonds. In such a case the bond is said to be non polar . However, if two atoms of different electronegativity share bonding electrons, the more electronegative atom will attract the bonding electrons becoming slightly negatively charged whereas the less electronegative atom assumes a slightly positive charge. Thus, the bond is said to be polar and the molecule has a dipole moment. For example the H-Br molecule is polar since Br is more electronegative than H, attracting the bonding electrons and giving H a partial positive charge and Br a partial negative charge.
The mere presence of a polar bonds does not guaranty a polar molecule. For example carbon dioxide O=C=O has two polar C=O bonds, yet the molecule is non polar because the geometry of CO₂ is linear.

Oxygen (2s², 2p⁴) undergoes also a hybridization , forming four sp³ hybrid orbitals. Each of the two lone electron pairs are occupying a hybrid orbital and the two single electrons occupy the two remaining hybrid orbitals. In case of water the bond angle between O and the two H is 104.5 degrees. This not exactly tetrahedral (109.3 degrees) as shown for CH₄ below, because the two lone electron pairs repell each other. The bent shape water causes a strong dipole moment. The 5 electrons in the outer shell of nitrogen in ammonia undergo also sp₃ hybridization, resulting in tetrahedral bond angles. The unbonded electron pair represents the negative side of the dipole whereas the three H-atoms are partially positively charged. Bond energies of polar bonds are relatively weak and range from 2 - 3 kcal/mol.

Hydrogen bonds
A hydrogen bond or H-bond is a special polar interaction between molecules which contain hydrogen atoms bonded to N, O or F. The H-bond of water in this example is formed through orbital overlap between the lone electron pair of oxygen and the partially positive charged hydrogen, bonded to another oxygen. The H-bond is strongest if the angle between O^...... H - O is 180^o. Average bond energies of H-bonds are 4 - 5 kcal/mol.

Hydrophobic Interactions
Hydrophobic interactions also called van der Waals forces are the weakest molecular interactions ( bond energy about 1 kcal/mol). Non polar bonds such as C-H can have temporary a dipole formation, which in turn can induce a temporary dipole in a neighboring molecule. This causes a weak attraction between non polar molecules.

Weak molecular interactions are extremely important in the association of biomolecules. For example the DNA double helix is held together largely by H-bonds between the nucleic acid bases. The specific association between enzymes and their substrates is brought about by weak molecular interactions such as ionic bonds, polar and hydrophobic interactions and H-bonds. Hormones form association complexes with their specific receptors through weak molecular interactions and thus initiate physiological responses. Because of their non polar structure biomembranes act as a barriers for polar molecules, which consequently need specific carrier molecules to be transported through the biomembrane. The importance of weak interactions is that they an can be form easily, causing a cellular response which is terminated as easily by the disruption of these forces.

Problems

1. Find the number of protons, neutrons and electrons in the following neutral atoms
   a) C-12
   b) C-14
   c) N-14
   d) O-18
   e) F-19
2. Calculate the atomic number and mass number of
   a) Mg
   b) Ca
   c) Cl-35
   d) C-13
   e) D
3. Explain the relationship between number of protons and electrons in an atom.
4. What is the name of the element with
   a) 9 protons and 10 neutrons
   b) 17 protons and 20 neutrons
   c) 15 protons and 17 neutrons
   d) 15 protons and 16 neutrons.
5. Oxygen has a mass of 2.66 x 10^-23 g. What is its atomic mass?
6. What are the differences between the most common types of nuclear radiations?
7. Tritium (H-3) emits beta radiation. What is the product of this nuclear reaction?
8. Po-214 is an alpha and gamma emitter. What is the product of this nuclear decay?
9. P-32 has a half life of 14 days. How many mg of P-32 are left after 56 days. The initial amount of P-32 was 120 mg
10. Sr-90 was one of the radioactive elements released into the atmosphere during the nuclear accident at Chernoble. Its half life = 29 years. How many years will it take before the radioactivity in a fall out area has decreased to 0.1 %?
11. How many electrons are allowed to be in a
    a) 1s orbital
    b) 2s orbital
    c) all 2p orbitals
    d) one single 4d orbital
    
12. Determine the electron configuration of
    a) N
    b) Si
    c) P-32
13. Identify the following elements:
    a) 1s², 2s², 2p¹ b) 1s², 2s², 2p⁵
    c) 1s², 2s², 2p⁶, 3s², 3p¹
    
14. Na and CL can form an ionic bond. Which atom becomes a cation and which atom becomes the anion?
15. which of the following structures are capable of H-bond formation
    
    

1. : P, N, E
   a) 6, 6, 6
   b) 6, 8, 6
   c) 7, 7, 7
   d) 8, 10, 8
   e) 9, 10, 9
2. : Atomic #, Mass #, # of neutrons
   a) 12, 24, 12
   b) 20, 40, 20
   c) 17, 35, 18
   d) 6, 13, 7
   e) 1, 2, 1
3. # of protons = # of electrons in an element (no net charge)
4. a) F, b) Cl-37, c) P-32, d) P-31
5. 16.02 amu - mass equivalent of energy lost in the formation of nucleus
6. alpha (He nuclei), beta (electrons), gamma radiation (short wave electromagnetic radiation)
7. He-3. Tritium is an isotope of hydrogen. It has a nucleus with one proton and 2 neutrons. Tritium is instable and its nucleus emits beta rediation. Beta radiation is the result of a neutron being converted into a proton and an electron. The electron escapes as beta radiation. Thus the remaining element is not Tritium anymore since the nucleus consist now of 2 protons and one neutron. The element with 2 protons (atomic # =2 ) is helium. Since it has only 1 neutron it is an isotope of helium : He-3 ( 3 is the mass number : 2 protons + 1 neutron) . Naturally occurring He has a mass number of 4 (2 protons + 2 neutrons )
8. Pb-210
   Po-214 has a mass number ( # of protons +neutrons) of 214. If Po-214 emits alpha radiation (= helium nuclei, He nucleus has 2 protons and 2 neutrons) the product of the decay is an element with 4 mass units less or to be specific, an element with 2 protons and 2 neutrons less than Po-214. In the periodic system Po has an atomic # = 84. The decay product has 2 protons less , therefore its atomic # = 82, which is lead (Pb). Its mass number must be 4 mass units less than Po-214 (since it looses a He nucleus) , thus Pb-210
9. 7.5 mg
10. 289 y
11. a) 2, b) 2, c) 6, d) 2
12. a)1s², 2s², 2p³
    b) 1s², 2s², 2p⁶, 3s², 3p²
    c) 1s², 2s², 2p⁶, 3s², 3p³,
13. a) B, b) F, c) Al
14. One electron from Na is transferred to CL
    Na ⁺ = cation, CL^- = anion
15. a), b), c), f) can form H-bonds
    d) the H-atom is not bonded to an oxygen or nitrogen atom and can not form an H-bond to the C=O group
    e) Even though a hydrogen bonded to an oxygen is in proximity to the oxygen of the C=O group, a H-bond can not be formed since the axis between O-H and O=C in not 180 degrees.