Electronic Structure of Matter PDF

Electronic Structure of Matter **ATOM** **smallest unit** into which matter can be divided without the release of electrically charged particles. It also is the smallest unit of matter that has the characteristic properties of a chemical. Each atom of every element consist of three basic subatomic...

Electronic Structure of Matter **ATOM** **smallest unit** into which matter can be divided without the release of electrically charged particles. It also is the smallest unit of matter that has the characteristic properties of a chemical. Each atom of every element consist of three basic subatomic particles: **electron**, **proton**, and **neutron.** **Three basic subatomic particles of an atom** **1.Electron** discovered by a physicist named **Joseph John Thompson ( 1856- 1940**).He discovered it through **cathode ray** experiment. Cathode rays are normally invisible beams of particles that can be contained within vacuum tubes for observation. If cathode rays are combined with fluorescent materials, they can be made visible to the naked eye. When he emitted the negatively charged particles or electrons into the cathode ray tube, these stream of electrons were drawn towards positively charged plate of the tube. **R.A Milikan** - another scientist who discovered the charge of electron to be 1.60 x10 ^-19^ coulomb **2. Proton** Is the **positively charge** sub -atomic particle that was proven of its existence through the gold foil experiment of **Ernest Rutherford** **3. Neutron** is the sub atomic particle that is electrically neutral. It was discovered by **James Chadwick** after he struck a very thin sheet of beryllium sheet released very high radiation similar to gamma rays **Comparison of the masses and charges of sub atomic particles**  Development of atomic Models **1. DEMOCRITUS** -(about 460BC) thought that all matter is composed of a very tiny invisible particles called **atomos**. These invisible particles are of the same particle having the same characteristics as the whole material 2\. **JOHN DALTON ( 1766- 1849 )** was eng.chemist who formulated the following a\. Elements are composed of very small particles called **atom** b\. Atoms of the same element are the same in size, shape, mass and chemical characteristics.The atoms of an element is exclusive to that particular element and is different from the atoms of another element. c\. Compounds are formed through the combination of elements combined in a fix proportion d\. Chemical reactions are a result of the arrangement of atoms and does not result to their destruction or creation. **3. Joseph John Thompson** -was a british physicist who based his model on a tradintional eng. dessert containing raisins. He presented the atom as a positive sphere( pudding) where electrons( Raisins) are embedded **4. Ernest Rutherford** -was a New Zealand born british who made the nuclear Model of the atom.He discovered that the atom consist of a positive dense center nucleus and proposed that electrons orbit the nucleus **Neils Bohr** -Emission Spectrum -Scientists observed that when materials are heated, they emit radiation that can be sensed by our naked eye. For ex. heated iron glows in specific way that can be seen by the naked eye.This glow is the visible spectrum emitted by the energized or heated iron atoms. -Another example of emission spectrum that can be observed when specific element heated is when fireworks explode, it actually exhibits various colors. these colors are the visible radiation emitted by the heated electrons of the elements **Bohr's Atomic Model - Hydrogen atom and atomic spectra** -he proposed of atomic model of a hydrogen atom to explain the electronic structure of the atom particularly the emission spectra of elements.He postulated that the electrons move in circular orbits around the nucleus. These orbits have specific energies.the orbit that closest to the nucleus has low energy and the orbit that is far from the nucleus has greater energy. when it is the 1st orbit (n1) it is in the ground state. When it is the excited state it gains extra energy and jump from ground state to higher orbit. This makes the atom unstable -The electron goes back to the ground state and releases a photon ( particle of light) which emits different visible spectra. This phenemenon was able to explain by Neils Bohr through his model of a hydrogen atom. \*Schrodinger's model is primarily mathematical ( based on mathematical solution to an equation that describes electrons in an atom \* The quantum mechanical model of atoms restricts the energy to certain values, like Bohr model and unlike the Bohr Model however it does not define the exact path an electron takes around the nucleus. Instead it estimates the probability of finding an electron in a certain region around the nucleus The probability of finding an electron in a certain region in space surrounding the nucleus can be illustrated as **fuzzy cloud**, sometimes referred to as an **electron cloud.** -It is most dense when the probability of finding an electron is high; less dense when the probability is low \- The probabilities of a fuzzy cloud are usually limited to the volume in which an electron is found 90% of the time. The electron probability cloud takes different shapes, depending upon the energy level of the electron, In quantum mechanical model, **these shapes** are called **atomic orbitals** **ATOMIC ORBITALS** The energy levels of electron are designated by the principal quantum numbers(n) The principal quantum numbers are refers as to the major principal energy level in an atom or the main energy level in an atom. The values assigned according to increasing energy may be n = 1,2,3,4,5,6,7 ans so on with 1 as the energy level closest to the nucleus. The average distance of an energy level from the nucleus increase as n increases. Electrons occupy energy sublevels within the principal or main energy level.Much as in school there are departments ex.. elementary and high school department ( principal or main energy level ) within each department there are different grade level or year level ( energy sublevel) **S - orbital are spherical** **p-orbitals - perpendicular to one another are shaped like dumbbell** **d- orbitals -four of the five kinds of this orbitals are shaped like clover leaves** **The f orbitals are more complex and harder to visualize** The principal quantum number is always equal to the number of sublevels in that principal energy level. The number of orbitals within a principal energy level is given by the formula N , which means that in the second main energy level n=2 there is a total of 2^2^ which one 2s orbital and three 2p orbitals. The maximum number of electrons that can occupy a principal energy level is given by the formula 2n^2^. So in the second principal energy level the maximum number of electrons will be 2(2) ^2^ or 8 electrons **Energy levels and sublevels and orbitals of electrons** **Electronic Configuration** the way in which electron are arranged around the nucleus of an atom it can be written in two ways: 1\. **using spdf notation** 2\. **orbital box notation** **for Hydrogen atom** principal energy level **1 s^1\ no.\ of\ electrons^** orbital **SPDF Notation** The electron above notation is located in \- 1st energy level \- s orbital **Orbital box Notation** -For Helium atom 1s^1^ The following are the rules to follow in writing electronic configuration of atom: **1. Pauli Exclusion Principle by wolfgang Pauli** \- only two electrons can occupy an orbital and their spins must be paired **2. Aufbau Principle- building up principle** \- Orbitals fill in order of increasing energy from lowest energy to highest energy **3. Hund's rule** \- When orbitals of equal energy are available but there are not enough electrons to fill all them, one electron is added to each orbital before a second electron is added to any one of them. **Mnemonic Device** - is used as guide in filling the sublevels with electrons. Use this when using the SPDF notation **QUANTUM MECHANICAL MODEL OF THE ATOM ( Electron Cloud)** The **quantum mechanical model of the atom** explains that the electrons surrounding the nucleus exist in the form of electron clouds called **orbitals**. These electrons can be found in energy levels and sublevels (subshells)consisting of orbitals where electrons reside. Important Concepts in the Quantum Mechanical Model of Atom ***[1. Main Energy Level]*** - various energy positions in the atom in which electrons may be located. They are referred to by the numbers 1.2.3.4.5.6.7 ***[2. Energy Sublevels]*** -Each energy is made up of alone or more sublevels in which electron of varying energies can reside. Sublevels are sometimes called energy subshells. An energy level can have s,p,d and f sublevels. ***[3. Electron Orbitals]*** - each energy sublevel is made of one or more electron orbitals in which electrons can be located within an atom. Orbitals are three dimensional electron cloud. The various sublevel are referred to as s orbitals, p orbitals, d orbitals and f orbitals. **3 Physicist led to the development of a better model of the atom** **1.Louie de Brogle 1924** - proposed that the electron (which is thought of as a particle) could also be thought as a wave. **2. Erwin Shrodinger**- use the wave idea to develop a mathematical equation to describe the hydrogen atom. He discovered that the electron move around the nucleus in a cloud not orbit.Orbital help us to predict the area where we find electron. The closer position to the nucleus the higher the chance to find electron **3. Werner Karl Heisenberg**- discovered that for a very small particle like electron its location cant be exactly be knownand how it is moving-this is called **UNcertainty Principle** ( you can't locate the exact position of an electron at any given time( too small too fast) **The electronic configuration of Elements** **Energy Level** **Sublevel** ------------------ -------------- 1 1s 2 2s,2p 3 3s,3p,3d 4 4s, 4p,4d,4f 5 5s,5p,5d,5f 6 6s,6p,6d 7 7s,7p **Sublevels** **No.of Orbitals** **Electron Capacity** --------------- -------------------- ----------------------- s 1 2 p 3 6 d 5 10 f 7 14 **Steps in writIng the electronic configuration of elements** 1\. Get the no.of electrons by using the atomic no.The atomic no. of neutral atom is also equal to the number of electrons ex Magnesium atomic no. is 12 2\. Distribute the electron in the different sublevels according to the order of filling.After filling a sub level in sequence and the rest of the sublevels until all the electrons are used.You can use electrons distribution mnemonic device to write electronic configuration ex. Magnesium atom atomic no is 12 Mg ^12^ - 1s^2^ 2s^2^ 2p^6^ 3s^2^ 3\. If an element is a positive ion (cation) or negative ion (anion), there will be a change in the number of electrons.A positive ion means that atom lost electrons and negative ion means a gain of electrons. 4\. Elements tend to follow the electronic configuration of the noble gas elements to be stable that is having 8 electrons in the s and p orbitals of their last energy level or 8 valence electron.When elements achieve the same electronic configuration with noble gas they are said to be **isoelectroic** Example:an argon atom and a (-1) chloride ion. Both have 18 electrons, arranged in the same numbers in the same shells. **PERIODIC TABLE** **Periodic table** \- is a system of arranging elements in the order of their increasing atomic weights and the recurrence of the properties **Antoine Lavoisier** \- father of the modern chemistry. He arranged the elements into groups of simple substances that do not compose by any means **JW Dobereiner** - in the early 1800s he formulated the law of triad in which he grouped the elements into sets of three. **John A.R Newlands** - he devised a law that explain that chemical groups repeated every eight elements and call it the Law of Octaves **Dmitri Ivanovich Mendeleev** - he created the first accepted version of the periodic table. he grouped the periodic table according to their atomic mass and as hed did he found that families had similar chemical properties **Henry Moseley** - was able to calculate the atomic no. of all known elements using experimental set up. He arranged the element into increasing atomic no. and formulated the periodic law which state that the properties of the elements are a periodic function of their atomic no. **Characteristics of Periodic Table** **Atomic number** - this refers to the no. of protons of an atom of an element. No two elements have the same no. of protons **Atomic mass** - refers to the weight of the atom. It is the sum of protons and neutrons of an atom **Symbols** - all elements have their own unique symbols.It can consist of a single capital letter or a capital letter and one or two lower case letters it is from the latin or english name of elements.It was John Jacob Berzelius who devised the present system of the chemical symbols **Valence Electron** - are the electrons in the outer energy level of an atom.These are the electrons that are transferred or shared when atoms bond together.In a period the no. of valence electrons increases from left to right **Metals** - are good conductors of heat and electricity, ductile, shiny and malleable **Non Metals** - are poor conductors of heat and electricity are not ductile, or malleable. solid non metals are breakable or brittle.they are dull. gases are non metals **Metalloids** - have properties of both metal and non metal **Families** - column of elements are called families or group. Elem. in each family have similar but not identical properties. All elem. in the family have the same no. of valence electrons **Periods**- Each horizontal row of elements is called a period. Element in period are not alike in properties. The first element in period is always an extremely active solid and the last element in a period is always an inactive gas **PERIODIC TRENDS** **Ionization Energy** is the energy required to removed an electron from an atom. TREND:It increases across a period, decreases down a group.The lower the ionization energy the easier it is to remove its valence electron. **Electron affinity** - this is the energy released by an atom when an electron is added to it. TREND: It increases across a period decreases down a group. An element that really accepts electron has higher electron affinity. These elements belongs to group Va-VII A **Electronegativity** - ability of an atom to attract or hold electrons. TREND: Its trend is the same with electeron affinity.It is a measure of the tendency of an atom to attract electrons.The higher the value the higher its tendency to attract electrons LEWIS STRUCTURE OF AN ATOM **Gilbert Newton Lewis** proposed that chemical bonds were formed between atoms as electrons from the atoms interacted with each other. Lewis had observed that many elements are most stable when they contained eight electrons in their valence shell. This observation gave rise to the octet rule. **Octet rule** states that atoms tend to lose, gain or share electrons in order to acquire a full set of valence electrons. Lewis Electron Dot Structure (LEDS), a shorthand to represent the valence electrons of an atom. The number is equal to the number of valence electron in representative elements (elements that belong to 1A-VIIA.) **LEDS** -consist of the chemical symbol for an element plus a dot for each valence electron Representative elements are elements in group IA- VIIIA. Always remember that the group number will be the valence electron which determines the number of dots present. **Group B elements** are the transitional metals **Group IA** - alkali metals ( metals forming compound with oxygen which when dissolved in water produce solutions that are strongly alkaline **Group II A** -alkaline earth metals( their compounds are not soluble in water but are found in ground deposits. **Type of Chemical Bond:** **IONIC BOND** [Compounds] are chemical substances made up of two or more elements that are chemically bound together in a fixed ratio. [Chemical bond] - refers to the strong attractive forces holding atoms together to form molecules and solids. **COVALENT BOND** **Covalent bond**- is another type of bond in which sharing of electrons takes place. Two Types : 1\. **Polar covalent bond** - the electronegativity difference is less than 1.9 and more than 0.4. 2\. **Non Polar Covalent Bond** - the electronegativity difference is equal to 0.4 or less to try if covalent bonding takes place.Apply the following steps; 1\. Get the total available valence electrons in a compound. example CO2 Carbon atom has 4 valence electrons Oxygen atom has 6 valence electrons total available Valence electrons (TAVE) = ( 1 C atom x 4) + ( 2 O atoms x 6 ) = 4 + 12 = 16 2\. Compute for the octet Rule requirement that each atom should have 8 valence electrons to become stable. Number of electrons based on Octet Rule = = ( 1 C atom x 8 ) + 2 O atoms x 8 = 8 + 16 =24 C. Subtract A from B , then divied the difference by 2 because a pair of shared electrons is equal to 1 bond. the quotient will give you the number of bonds around the central atom No. of bonds = [ ( 24-16) ] 2 = 4 Thus there will be 4 bonds surrounding a carbon atom as shown in Lewis Structure **CHEMICAL NOMENCLATURE** Chemical Nomenclature \- the naming of substances From latin word NOMEn ( name) and Calare ( to call) Steps: 1\. write the symbols for each element 2\. On the upper right hand, write the oxidation number above each element 3\. Cross oxidation number of one element to the other element. The oxidation number will be the subscript of the other element ( take the absolute value) 4\. Rewrite the formula and drop the charges 5\. Reduce the subscript to their simplest form if needed. 6\. The sum of oxidation number must be equal to zero **NAMING OF COMPOUNDS** BINARY COMPOUNDS: Rules: Name the metallic element firs followed by non metallic elements, then end with - ide. example: Al~2~O~3\ -~ Aluminum Oxide KBr - Potassium bromide Non metal and non metal combination Rule: mention the prefix of the first element then the name of the first element followed by the prefix of the second element, then the name of the second element, ending in -ide GREEK PREFIXES: 1 - mono 6 - hexa 2 - di 7 - hepta 3 - tri 8 - octa 4 - tetra 9- nona 5 - penta 10 - deca example: N~2~O~5~ = di + nitrogen + penta + oxygen = Dinitrogen pentoxide CO~2~ = Mono + carbon + di + oxygen = monocarbon dioxide or carbon dioxide TERNARY COMPOUNDS \***POLYATOMIC IONS** RULE: Name first element followed by the name of the name of the polyatomic ion attached into it. examples: Mg (NO3)2 - Magnesium Nitrate \* (NO3)2 is the polyatomic ion TERNARY COMPOUNDS **\* Metal with variable oxidatiopn no. and non metals combinations** **Stock System** - indicates the metal ionic charge by writing it in Roman Numerals within a parenthesis example: Ion Stock system Fe+2 Iron ( II) Fe+3 Iron ( III) **Classical Method** - uses the latin name of the metal ,if the metal uses its highest oxidation number , rootword of the latin name ofd the metal followed by - ic, if it uses lowest oxidation number root word of the latin name of the metal followed by - ous ion Latin Name Stock system ------ ------------ -------------- Fe+2 Ferrum Ferrous Fe+3 Ferrum Ferric Cu+1 Cuprum Cuprous Cu+2 Cuprum Cupric **Structure of Carbon Atoms** **The Hydrocarbons** - are organic compounds that contain only carbon and hydrogen atoms - This ability of carbon to form big molecules has enabled nature to form living matter from simple cells to complex tissues. - Hydrocarbons such as methane, ethane and butane are components of natural gas - Hydrocarbons are grouped into families namely:alkanes, alkenes and alkynes.The compound from each group have certain structures that make their properties different from each other. **A. ALKANES** - are hydrocarbons with only single bonds between carbon atoms in the compound - are considered sataturated hydrocarbons because all the valence electrons of each carbon atom are already bonded.This means that alkanes can no longer bond with other additional elements. **Example:** M**ethane** **Phase** -gas **Condensed structural formula:** **CH~4~** - **Structural formula** **H** **H - C - H** **H** **Boiling point** = **-162** **Alkenes** - are hydrocarbons that contain a double bond between carbon atoms - are considered unsaturated hydrocarbons because bonds with another carbon or other elements can still be formed from double bonds - Example: **ETHENE** - **Phase = Gas** - **Condensed structural formula : CH2=CH2** - **Structural Formula :** - **H H** - **C = C** - **H H** - **boiling Point: -142** **ALKYNES** - are hydrocarbons compounds with a tripple bond between two carbon atoms. - they are also unsaturated hydrocarbon compounds Example : **ETHYNE** **Phase : Gas** **Condensed structural formula : C~2~CH~2~** **Structural Formula** **H- C C- H** **Boiling point : -84** **Two major classes of organic compounds** **1. Hydrocarbons** -contains only 1 and hydrogen in the structure 2\. **Substituted Hydrocarbons** - are hydrocarbons in which one or more hydrogen hydrogen atom is/are replaced by another element **Naming and Writing formulas of hydrocarbons** **Molecular formula** - gives the kind and number of each type of atoms composing an organic compound.This type of formula , however does not show the actual bonding pattern and atom arrangements of the compound.That is why there is a possibility that two or more compounds may have the same molecular formulas example: the molecular formula of propene is **C3H6** **Structural formula** - is a formula that shows the actual structure of the molecule and gives **all the atoms in a sequential order indicating bonding patterns** 1\. Expanded 2\. Condensed **Example Ethene** **CH2=CH2 - condensed** **H H** **C = - Expanded** **H H** **Guideline for naming hydrocarbons and otherorganic compounds** INTERNATIONAL UNION OF PURE AND APPLIED CHEMISTRY (IUPAC) - ALKANES NOMENCLATURE 1\. Find and name the longest continuous carbon chain 2\. number the longest continuous chain carbon atoms starting from the end that gives the lowest numbers to the branches on the chain 3\. indicate the postion of the alkylgroup/s by name and number 4\. if there is more than one of the same substituent group on the chain ,indicate the number by a prefix : di - 2,tri- 3, tetra - 4 and so on..Considered another methyl group added **Greek prefixes** 1 -mono 6 - hexa 2- di 7- hepta 3- tri 8- octa 4- tetra 9- nona 5- penta 10 deca  Nomenclature for Alkenes and alkynes 1\. Select the longest carbon carbon chain that contains the double or tripple bond 2\. Names this parent compound as alkene ( double bond) or alkyns for tripple bond 3.Number the carbon chain of the parent compound starting with the end nearer to the double or tripple bond.Use the smaller of the two nos. on the double or tripple bonded carbon atoms to indicate the position of the double or tripple bond.Place this no. infront of the alkens or alkyn name 4\. side chains and the other groups are treated as in naming alkanes, by prefixing a number to them corresponding to the carbon atom to which they are bonded