Introduction to Biochemistry for Nurses BSN Post RN

Basic Concepts in Chemistry for Nurses 

Brief introduction to basic concepts of chemistry before biochemistry for nurses.

Chemistry

    The branch of science which deals with the composition, structure, properties and reactions of matter is called chemistry.

Biochemistry

    It is the branch of chemistry in which we study the structure, composition, and chemical reactions of substances found in living organisms. It covers all chemical processes taking place in living organisms, such as synthesis and metabolism of bio molecules like carbohydrates, proteins and fats. Biochemistry emerged as a separate discipline when scientists began to study how living things obtain energy from food or how the fundamental biological changes occur during a disease. Examples of applications of biochemistry are in the fields of medicine, food science and agriculture, etc.

Matter

    Matter is simply defined as anything that has mass and occupies space. Our bodies as well as all the things around us are examples of matter. In chemistry, we study all types of matters that can exist in any of three physical states: solid, liquid or gas.

Substance

    A piece of matter in pure form is termed as a substance. Every substance has a fixed composition and specific properties or characteristics.

Mixture

    Impure matter is called a mixture; which can be homogeneous or heterogeneous in its composition.

Physical properties

    The properties those are associated with the physical state of the substance are called physical properties like color, smell, taste, hardness, shape of crystal, solubility, melting or boiling points, etc. For example, when ice is heated, it melts to form water. When water is further heated, it boils to give steam. In this entire process only the physical states of water change whereas its chemical composition remains the same.

 Chemical properties

    The chemical properties depend upon the composition of the substance. When a substance undergoes a chemical change, its composition changes and new substances are formed. For example, decomposition of water is a chemical change as it produces hydrogen and oxygen gases.

Elements

    It is a substance made up of same type of atoms, having same atomic number and cannot be decomposed into simple substances by ordinary chemical means. Elements may be solids, liquids or gases. Majority of the elements exist as solids e.g., sodium, copper, zinc, gold, etc. There are very few elements which occur in liquid state e.g., mercury and bromine. A few elements exist as gases e.g., nitrogen, oxygen, chlorine and hydrogen.

    On the basis of their properties, elements are divided into metals, non-metals and metalloids. About 80 percent of the elements are metals.

Symbols

    Elements are represented by symbols, which are abbreviations for the name of elements. A symbol is taken from the name of that element in English, Latin, Greek or German. If it is one letter, it will be capital as H for Hydrogen, N for Nitrogen and C for Carbon etc. In case of two letters symbol, only first letter is capital e.g., Ca for Calcium, Na for Sodium and Cl for Chlorine.

Valency

    The unique property of an element is valency. It is combining capacity of an element with other elements. It depends upon the number of electrons in the outermost shell.

Some Elements and Radicals with their Symbols and Common Valences

Element/Radical	Symbol	Valency	Element/Radical	Symbol	Valency
Sodium	Na	1	Hydrogen	H	1
Silver	Ag	1	Chlorine	CI	1
Magnesium	Mg	2	Bromine	Br	1
Calcium	Ca	2	Iodine	I	1
Barium	Ba	2	Oxygen	O	2
Zinc	Zn	2	Sulphur	S	2
Copper	Cu	1,2	Nitrogen	N	3
Mercury	Hg	1,2	Phosphorus	P	3,5
Iron	Fe	2,3	Boron	B	3
Aluminum	Al	3	Arsenic	As	3
Chromium	Cr	3	Carbon	C	4
Ammonium	NH4+	1	Carbonate	CO32-	2
Hydronium	H3O+	1	Sulphate	SO42-	2
Hydroxide	OH-	1	Sulphite	SO32-	2
Cyanide	CN-	1	Thiosulphate	S2O32-	2
Bisulphate	HSO4-	1	Nitride	N3-	3
Bicarbonate	HCO3-	1	Phosphate	PO43-	3

 Compound

    Compound is a substance made up of two or more elements chemically combined together in a fixed ratio by mass.

    For example, carbon dioxide is formed when elements of carbon and oxygen combine chemically in a fixed ratio of 12:32 or 3:8 by mass. Similarly, water is a compound formed by a chemical combination between hydrogen and oxygen in a fixed ratio of 1:8 by mass.

    Compounds can be classified as ionic or covalent. Ionic compounds do not exist in independent molecular form. They form a three-dimensional crystal lattice, in which each ion is surrounded by oppositely charged ions. These oppositely charged ions attract each other very strongly, as a result ionic compounds have high melting and boiling points. These compounds are represented by formula units e.g., NaCl, KBr, CuSO_4.

Molecular formula

    The covalent compounds mostly exist in molecular form. A molecule is a true representative of the covalent compound and its formula is called molecular formula e.g. H₂O, HC1, H₂SO₄, Ch₄etc.

Some Common Compounds with their Formulae

Compound	Chemical Formula
Water	H₂O
Sodium chloride (Common salt)	NaCl
Silicon dioxide (Sand)	SiO2
Sodium hydroxide (Caustic Soda)	NaOH
Sodium carbonate (Washing Soda)	Na2CO3.10H₂O
Calcium oxide (Quick Lime)	CaO
Calcium carbonate (Limestone)	CaCO3
Sugar	C12H22O11
Sulphuric acid	H2SO4
Ammonia	NH3

Mixture

    When two or more elements or compounds mix up physically without any fixed ratio, they form a mixture. The mixture can be separated into parent components by physical methods such as distillation, filtration, evaporation, crystallization or magnetization.

Homogeneous mixtures

    Mixtures that have uniform composition throughout are called homogeneous mixtures e.g., air, gasoline, ice cream.

Heterogeneous mixtures

    Heterogeneous mixtures are those in which composition is not uniform throughout e.g., soil, rock and wood.

Difference between a Compound and a Mixture

Compound	Chemical Formula
1.     It is formed by a chemical combination of atoms of the elements	1.     Mixture is formed by the simple mixing up of the substances.
2.     The constituents lose their identity and form a new substance having entirely different properties from them.	2.     Mixture shows the properties of the constituents.
3.     Compounds always have fixed composition by mass.	3.     Mixtures do not have fixed composition.
4.     The components cannot be separated by physical means.	4.     The components can be separated by simple physical methods.
5.     Every compound is represented by a chemical formula.	5.     It consists of two or more components and does not have any chemical formula.
6.     Compounds have homogenous composition.	6.     They may be homogeneous or heterogeneous in composition
7.     Compounds have sharp and fixed melting points.	7.     Mixtures do not have sharp and fixed melting points.

Atomic Number and Mass Number

    The atomic number of an element is equal to the number of protons present in the nucleus of its atoms. It is represented by symbol ‘Z’. As all atoms of an element have the same number of protons in their nuclei, they have the same atomic number.

    Hence, each element has a specific atomic number termed as its identification number. For example, all hydrogen atoms have 1 proton, their atomic number is Z=l. All atoms in carbon have 6 protons, their atomic number is Z=6. Similarly, in oxygen all atoms have 8 protons having atomic number Z=8 and Sulphur having 16 protons shows atomic number Z = 16.

    The mass number is the sum of number of protons and neutrons present in the nucleus of an atom. It is represented by symbol 'A'.

    It is calculated as A=Z+n where n is the number of neutrons.

     Each proton and neutron have 1 amu mass. For example, hydrogen atom has one proton and no neutron in its nucleus, its mass number A=l+0 =1. Carbon atom has 6 protons and 6 neutrons, hence its mass number A=12. Atomic numbers and mass numbers of a few elements are given in Table.

Some Elements along with their Atomic and Mass Numbers

Electrons	Number of Protons	Number of Neutrons	Atomic Number Z	Mass Number
Hydrogen	1	0	1	1
Carbon	6	6	6	12
Nitrogen	7	7	7	14
Oxygen	8	8	8	16
Fluorine	9	10	9	19
Sodium	11	12	11	23
Magnesium	12	12	12	24
Potassium	19	20	19	39
Calcium	20	20	20	40

Relative Atomic Mass and Atomic Mass Unit

    The relative atomic mass of an element is the average mass of the atoms of that element as compared to 1/12 (one-twelfth) the mass of an atom of carbon-12 isotope (an element having different mass number but same atomic number).

    The unit for relative atomic masses is called atomic mass unit, with symbol 'amu'. One atomic mass the unit is 1/12 the mass of one atom of carbon-12. When this atomic mass unit is expressed in grams, it is:

How to write a Chemical Formula

    Compounds are represented by chemical formulae as elements are represented by symbols. Chemical formulae of compounds are written keeping the following steps in consideration:

·        Symbols of two elements are written side by side, in the order of positive ion first and negative ion later.

·        The valency of each ion is written on the right top corner of its + 2+ -- symbol, e.g., Na, Ca, CI and O₂.

·        This valency of each ion is brought to the lower right corner of other ion by 'cross exchange' method, e.g.

They are written as:

·            If the valences are same, they are offset and are not written in the chemical formula. But if they are different, they are indicated as such at the same position, e.g., in case of sodium chloride both the valences are offset and formula is written as NaCl, whereas, calcium chloride is represented by formula CaCl_2.

·            If an ion is a combination of two or more atoms which is called radical, bearing a net charge on it, e.g., SO₄^2-(sulphate) and PO (phosphate), then the net charge represents the valency of the radical. The chemical formula of such compounds is written as explained in (iii) and (iv); writing the negative radical within the parenthesis. For example, chemical formula of aluminum sulphate is written as Al₂ (SO₄ )₃and that of calcium phosphate as Ca₃ (PO₄)_2.

Theories And Experiments Related to Structure of Atom

    According to Dalton, an atom is an indivisible, hard, dense sphere. Atoms of the same element are alike. They combine in different ways to form compounds. In the light of Dalton's atomic theory, scientists performed a series of experiments. But in the late 1800's and early 1900's, scientists discovered new subatomic particles.

    In 1886, Goldstein discovered positively charged particles called protons. In 1897, J.J. Thomson found in an atom, the negatively charged particles known as electrons. It was established that electrons and protons are fundamental particles of matter. Based upon these observations Thomson put forth his “plum pudding” theory. He postulated that atoms were solid structures of positively charge with tiny negative particles stuck inside. It is like plums in the pudding.

Atom

    A particle of matter that uniquely defines a chemical element. An atom consists of a central nucleus that is surrounded by one or more negatively charged electrons. The nucleus is positively charged and contains one or more relatively heavy particles known as protons and neutrons.

Ion

    An ion is an atom or group of atoms that has an electric charge. Ions with a positive charge are called cations. Ions with a negative charge are called anions. To form an ion, an element must gain or lose an electron. Gaining electrons or losing electrons creates an ion. 

If an atom gains an electron, it has more electrons than protons, creating an overall negatively charged atom of an element.Many normal substances exist in the body as ions. Common examples include sodium, potassium, calcium, chloride, and bicarbonate.

Radical

    Also known as a free radical, is an atom, molecule, or ion that has at least one unpaired valence electron. With some exceptions, these unpaired electrons make radicals highly chemically reactive.

Periodic Table

    On the basis of this law, the elements known at that time, were arranged in the form of a table which is known as periodic table.

Dobereiner's Triads

    A German chemist Dobereiner observed relationship between atomic masses of several groups of three elements called triads. In these groups, the central or middle element had atomic mass average of the other two elements.

Newlands Octaves

    After successful determination of correct atomic masses of elements by Cannizzaro in 1860, attempts were again initiated to organize elements. In 1864 British chemist Newlands put forward his observations in the form of 'law of octaves'.

Mendeleev's Periodic Table

    Russian chemist, Mendeleev arranged the known elements (only 63) in order of increasing atomic masses, in horizontal rows called periods. So that elements with similar properties were in the same vertical columns. This arrangement of elements was called Periodic Table. He put forward the results of his work in the form of periodic law, which is stated as "properties of the elements are periodic functions of their atomic masses".

Modern Periodic Table

    The modern periodic table is based upon the arrangement of elements according to increasing atomic number. When the elements are arranged according to increasing atomic number from left to right in a horizontal row, properties of elements were found repeating after regular intervals such that elements of similar properties and similar configuration are placed in the same group.

    It was observed that after every eighth element, ninth element had similar properties to the first element. For example, sodium (Z=11) had similar properties to lithium (Z=3). After atomic number 18, every nineteenth element was showing similar behavior. So, the long rows of elements were cut into rows of eight and eighteen elements and placed one above the other so that a table of vertical and horizontal rows was obtained.

Long form of Periodic Table

    The significance of atomic number in the arrangement of elements in the modern periodic table lies in the fact that as electronic configuration is based upon atomic number, so the arrangement of elements according to increasing atomic number shows the periodicity (repetition of properties after regular intervals) in the electronic configuration of the elements that leads to periodicity in their properties. Hence, the arrangement of elements based on their electronic configuration created a long form of periodic

    The horizontal rows of elements in the periodic table are called periods. The elements in a period have continuously increasing atomic number i.e., continuously changing electronic configuration along a period.

    The vertical columns in the periodic table are called groups. These groups are numbered from left to right as 1 to 18. The elements in a group do not have continuously increasing atomic numbers.

Salient Features of Long Form of Periodic Table:

i.                   This table consists of seven horizontal rows called periods.

ii.                 First period consists of only two elements. Second and third periods consist of 8 elements each. Fourth and fifth periods consist of 18 elements each. Sixth period has 32 elements while seventh period has 23 elements and is incomplete.

iii.              Elements of a period show different properties.

iv.              There are 18 vertical columns in the periodic table numbered 1 to 18 from left to right, which are called groups.

v.                 The elements of a group show similar chemical properties.

vi.              Elements are classified into four blocks depending upon the type of the subshell which gets the last electron.

Periods

    First period is called short period. It consists of only two elements, hydrogen and helium. Second and third periods are called normal periods. Each of them has eight elements in it. Second period consists of lithium, beryllium, boron, carbon, nitrogen, oxygen, fluorine and ends at neon, a noble gas. Fourth and fifth periods are called long periods. Each one of them consists of eighteen elements. Whereas, sixth and seventh periods are called very long periods.

     In these periods after atomic number 57 and 89, two series of fourteen elements each, were accommodated. Because of space problem, these two series were placed separately below the normal periodic table to keep it in a manageable and presentable form. Since the two series start after Lanthanum (Z=57) and Actinium (Z=89), so these two series of elements are named as Lanthanides and Actinides respectively. Table 3.1 shows the distribution of elements in periods.

     All the periods except the first period start with an alkali metal and end at a noble gas. It is to be observed that number of elements in a period is fixed because of maximum number of electrons which can be accommodated in the particular valence shell of the elements.

Groups

 Group 1 consists of hydrogen, lithium, sodium, potassium, rubidium, cesium and francium. Although elements of a group do not have continuously increasing atomic numbers, yet they have similar electronic configuration in their valence shells. That is the reason elements of a group are also called a family. For example, all the group 1 elements have one electron in their valence shells, they are given the family name of alkali metals. 

    The groups 1 and 2 and 13 to 17 contain the normal elements. In the normal elements, all the inner shells are completely filled with electrons, only the outermost shells are incomplete. For example, group 17 elements (halogens) have 7 electrons in their valence (outermost) shell. The groups 3 to 12 are called transition elements. In these elements 'a f ' sub-shell is in the process of completion. 

Shells 

    In Bohr's Atomic model electrons revolve around the nucleus in a specific circular path known as orbit or called a shell. The shells have stationary energy levels; the energy of each shell is constant. Every stationary orbit or shell is associated with a definite amount of energy.There are  K, L M ,N shells within an atom.

Sub shell 

    The grouping of electrons in a shell according to the shape of the region of space they occupy. Within each sub shell, electrons are grouped into orbitals, regions of space within an atom where the specific electrons are most likely to be found.There are s, p, d, f sub shells. 

Shell/Orbit  (Capital Letters)	No of Electrons	Sub Shell/Orbitals  (Small Letters)	No of Electrons
K	2	s	2
L	8	s,p	2,6   (2+6=8)
M	18	s,p,d	2,6,10   (2+6=10=18)
N	32	s,p,d,f,	2,6,10,14  (2+6=10=14=32)

Blocks 

    Block of the periodic table of elements is a set of adjacent groups of elements. The respective highest energy electrons in each element in a block belong to the same atomic orbital type. There are  4 blocks s, p, d, and f according to sub shell or orbital.

    A block of the periodic table is represent their atomic orbitals their valence electrons or vacancies lie in. 

Periodicity OF Properties

Atomic Size and Atomic Radius

    Half of the distance between the nuclei of the two bonded atoms is referred as the atomic radius of the atom. For example, the distance between the nuclei of two carbon atoms in its elemental form is 154 pm, it means its half 77 pm is radius of carbon atom.

    Half of the distance between the nuclei of the two bonded atoms is referred as the atomic radius of the atom. For example, the distance between the nuclei of two carbon atoms in its elemental form is 154 pm, it means its half 77 pm is radius of carbon atotomic number, the effective nuclear charge increases gradually because of addition of more and more protons in the nucleus. 

    But on the other hand, addition of electrons takes place in the same valence shell i.e., shells do not increase. There is gradual increase of effective nuclear charge which increases due to addition of protons. This force pulls down or contracts the outermost shell towards the nucleus. For example, atomic size in period 2 decreases from Li (152 pm) to Ne (69) pm.

    The size of atoms or their radii increases from top to bottom in a group. It is because a new shell of electrons is added up in the successive period, which decreases the effective nuclear charge. The trend of atomic size of transition elements has slight variation when we consider this series in a period. The atomic size of the elements first reduces or atom contracts and then there is increase in it when we move from left to right in 4th period.

Shielding Effect

    Valance electron experiences less nuclear charge than that of the actual charge, which is called effective nuclear charge (Z ). It means that the eff electrons present in the inner shells screen or shield the force of attraction of nucleus felt by the valence shell electrons. This is called shielding effect. With increase of atomic number, the number of electrons in an atom also increases, that results in increase of shielding effect.

    Valance electron experiences less nuclear charge than that of the actual charge, which is called effective nuclear charge (Z ). It means that the eff electrons present in the inner shells screen or shield the force of attraction of nucleus felt by the valence shell electrons. This is called shielding effect. With increase of atomic number, the number of electrons in an atom also increases, that results in increase of shielding effect.

Ionization Energy

     The ionization energy is the amount of energy required to remove the most loosely bound electron from the valence shell of an isolated gaseous atom. The amount of energy needed to remove successive electrons present in an atom increases. If there is only 1 electron in the valence shell, the energy required to remove it will be called first ionization energy. For example, the first ionization energy of sodium atom -1 is + 496 kJmol.

    If we move from left to right in a period, the value of ionization energy increases. It is because the size of atoms reduces and valence electrons are held strongly by the electrostatic force of nucleus. Therefore, elements on left side of the periodic table have low ionization energies as compared to those on right side of the periodic table as shown for the 2nd period.

    As we move down the group more and more shells lie between the valence shell and the nucleus of the atom, these additional shells reduce the electrostatic force felt by the electrons present in the outermost shell. Resultantly the valence shell electrons can be taken away easily. Therefore, ionization energy of elements decreases from top to bottom in a group.

Electron Affinity

     Electron Affinity is defined as the amount of energy released when an electron is added in the outermost shell of an isolated gaseous atom.

    Affinity means attraction. Therefore, electron affinity means tendency of an atom to accept an electron to form an anion. For example, the electron affinity of fluorine -1 is -328 kJ mol i.e., one mole atom of fluorine release 328 kJ of energy to form one mole of fluoride ions. Let us discuss the trend of electron affinity in the periodic table. Electron affinity values increase from left to right in the period.

    The reason for this increase is, as the size of atoms decreases in a period, the attraction of the nucleus for the incoming electron increases. That means more is attraction for the electron, more energy will be released. In group electron affinity values decrease from top to bottom because the size of atoms increases down the group. With the increase in size of atom shielding effect increases that results in poor attraction for the incoming electron i.e., less energy is released out. For example, as the size of iodine atom is bigger than chlorine, its electron affinity is less than iodine, as given in the adjacent table.

 Electronegativity

     The ability of an atom to attract the shared pair of electrons towards itself in a molecule, is called electronegativity.

    The trend of electronegativity is same as of ionization energy and electron affinity. It increases in a period from left to right because higher Z shortens distance eff from the nucleus of the shared pair of electrons. This enhances the power to attract the shared pair of electrons. For example, electronegativity values of group 2 are as follow:

    It generally decreases down a group because size of the atom increases. Thus, attraction for the shared pair of electrons weakens. For example, electronegativity values of group 17 (halogens) are presented here.

Concentration Units

    Concentration is the proportion of a solute in a solution. It is also a ratio of the amount of solute to the amount of solution or ratio of amount of solute to the amount of the solvent.

Percentage

    Percentage unit of concentration refers to the percentage of solute present in a solution. The percentage of solute can be expressed by mass or by volume. It can be expressed in terms of percentage composition by four different ways

Percentage - mass/mass (%m/m)

    It is the number of grams of solute in 100 grams of solution. For example, 10% m/m sugar solution means that 10 g of sugar is dissolved in 90 g of water to make 100 g of solution. Calculation of this ratio is carried out by using the following formula:

Percentage - mass/volume (%m/v)

    It is the number of grams of solute dissolved in 100 cm (parts by volume) of the 3 solutions. For example, 10 % m/v sugar solution contains 10 g of sugar in 100 cm of the solution. The exact volume of solvent is not mentioned or it is not known.

Percentage - volume/mass (%v/m)

    It is the volume in cm of a solute dissolved in 100 g of the solution. For example, 10 % v/m alcohol solution in water means 10 cm of alcohol is dissolved in (unknown) volume of water so that the total mass of the solution is 100 g. In such solutions the mass of solution is under consideration, total volume of the solution is not considered.

Percentage - volume /volume (% v/v)

    It is the volume in cm of a solute dissolved per 100 cm of the solution. For example, 30 percent alcohol solution means 30 cm of alcohol dissolved in sufficient amount of water, so that the total volume of the solution becomes 100 cm³

Chemical formula

    The chemical formula of a compound is a symbolic representation of its chemical composition. Chemical formulae provide insight into the elements that constitute the molecules of a compound and also the ratio in which the atoms of these elements combine to form such molecules.

Chemical formula for Aluminum sulfate:   

Chemical reactions 

    Chemical reactions occur when chemical bonds between atoms are formed or broken. The substances that go into a chemical reaction are called the reactants, and the substances produced at the end of the reaction are known as the products.

For example, the reaction for breakdown of hydrogen (H₂O₂) peroxide into water and oxygen can be written as:

Equations:

2 H₂O₂ (Hydrogen Peroxide) →2 H₂O₂  (Water)  + O₂(oxygen)

      2H₂+O₂→2 H₂O

Acids and bases are recognized by their characteristic properties, such as:

Acids	Bases
Acids have sour taste. For example, unripe citrus fruits or lemon juice.	Bases have bitter taste and feel slippery,for example, soap is slippery to touch.
They turn blue litmus red.	They turn red litmus blue.
They are corrosive in concentrated form.	They are non-corrosive except concentrated forms of NaOH and KOH.

Arrhenius Concept of Acids and Bases

    Acid is a substance which dissociates in aqueous solution to give hydrogen ions. For example, substances such as HC1, HNO₃, CH3 COOH, HCN, etc., are acids because they ionize in aqueous solutions to provide H+ ions.

Base is a substance which dissociates in aqueous solution to give hydroxide ions.The substances such as NaOH, KOH, NH4 OH, Ca (OH)₂ etc. are bases because these compoundsionize in aqueous solutions to provide OH ions.

Redox reactions,Oxidation,Reduction

    Redox reactions are oxidation-reduction chemical reactions in which the reactants undergo a change in their oxidation states. The term ‘redox’ is a short form of reduction-oxidation. All the redox reactions can be broken down into two different processes: a reduction process and an oxidation process.

    A redox reaction can be defined as a chemical reaction in which electrons are transferred between two reactants participating in it. This transfer of electrons can be identified by observing the changes in the oxidation states of the reacting species.

    The loss of electrons and the corresponding increase in the oxidation state of a given reactant is called oxidation. The gain of electrons and the corresponding decrease in the oxidation state of a reactant is called reduction.

• 2NaH → 2Na + H₂
• 2H₂O → 2H₂ + O₂
• Na₂CO₃ → Na₂O + CO₂

Examples of Redox Reactions

    A few examples of redox reactions, along with their oxidation and reduction half-reactions, are provided in this subsection.

Reaction between Hydrogen and Fluorine

    In the reaction between hydrogen and fluorine, the hydrogen is oxidized, whereas the fluorine is reduced. The reaction can be written as follows.

H₂ + F₂ → 2HF

The oxidation half-reaction is: H₂ → 2H⁺ + 2e^–

The reduction half-reaction is: F₂ + 2e^– → 2F^–

The hydrogen and fluorine ions go on to combine in order to form hydrogen fluoride.

Chemical bonding

    Any of the interactions that account for the association of atoms into molecules, ions, crystals, and other stable species that make up the familiar substances of the everyday world. When atoms approach one another, their nuclei and electrons interact and tend to distribute themselves in space in such a way that the total energy is lower than it would be in any alternative arrangement. If the total energy of a group of atoms is lower than the sum of the energies of the component atoms, they then bond together and the energy lowering is the bonding energy.

 Ionic Bond

    As the name suggests, ionic bonds are the result of attraction between ions. Ions are formed when an atom loses or gains an electron. These types of bonds are usually formed between a metal and a nonmetal.

    Sodium (Na) and chlorine (Cl) combine to form stable crystals of sodium chloride (NaCl), also known as table salt.

        Magnesium (Mg) and oxygen (O) combine to form magnesium oxide (MgO).

Potassium (K) and chlorine (Cl) combine to form potassium chloride (KCl)

Calcium (Ca) and fluoride (F) combine to form calcium fluoride (CaF₂)

Covalent bond

    In a covalent bond, an atom shares one or more pairs of electrons with another atom and forms a bond. This exchange of electrons occurs because the atoms must obey the octet rule (noble gas configuration) when bonding. This type of bonding is common between two nonmetals. The covalent bond is the strongest and most common form of chemical bonding in living organisms. Together with the ionic bond, they form the two most important chemical bonds.

    A covalent bond can be divided into a non-polar covalent bond and a polar covalent bond. In a non-polar covalent bond, the electrons are shared equally between the two atoms. In contrast, in polar covalent bonds, the electrons are unevenly distributed between the atoms.

Examples

Two iodine atoms (I) combine to form iodine gas (I₂).

A carbon atom (C) combines with two oxygen atoms (O) and forms a double covalent bond in carbon dioxide (CO₂).

Two hydrogen atoms (H) combine with an oxygen atom (O) to form a polar water molecule (H₂O).

Boron (B) and three hydrogen atoms (H) combine to form the polar borane (BH₃).

Metallic Bonds

    A metallic bond is a force that holds atoms in a metallic substance together. Such a solid consists of densely packed atoms, with the outermost electron shell of each metal atom overlapping with a large number of neighboring atoms. As a result, valence electrons move freely from one atom to another. They are not assigned to any particular pair of atoms. This behavior is called non-localization.

Examples

Metallic sodium

Frustrate

Copper cable

Other types of chemical bonds
Vander Waals Forces/Bond

    Neutral molecules are held together by weak electrical forces known as van der Waals forces. The van der Waals force is a general term used to define the attraction of intermolecular forces between molecules. This type of chemical bond is the weakest of all bonds.

Examples include hydrogen bonds, London dispersion forces, and dipole-dipole forces.

Dipole Dipole Forces 

    The attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. Dipole-dipole forces have strengths that range from 5 kJ to 20 kJ per mole.

Dipole Induce Dipole

    The  weak attraction that results when a polar molecule induces a dipole in an atom or in a non polar molecule by disturbing the arrangement of electrons in the non polar species.

Hydrogen Bonding

    A hydrogen bond is a chemical bond between a hydrogen atom and an electronegative atom. However, this is not an ionic or covalent bond, but a special type of dipole-dipole attraction between molecules. First, the hydrogen atom covalently bonds to a strongly electronegative atom, resulting in a positive charge, which is then attracted to an electronegative atom, resulting in hydrogen bonding.

Examples

The hydrogen atom of a water molecule combines with the oxygen atom of another molecule. This connection is of great importance in the ice.

In chloroform (CH₃Cl) and ammonia (NH₃), hydrogen bonding occurs between the hydrogen of one molecule and the carbon/nitrogen of another.

Nitrogen bases present in DNA are held together by hydrogen bonding.

Peptide bond

    Within a protein, several amino acids are linked together by peptide bonds, forming a long chain. Peptide bonds are formed by a biochemical reaction that extracts a water molecule when it joins the amino group of one amino acid to the carboxyl group of adjacent amino acids. In addition to peptide bonds, hydrogen bonds, ionic bonds and disulfide bonds are also common in proteins.

Examples include polypeptides such as insulin and growth hormone.