Name: ……………………………………………… Class:………………………………………………… Number:…………………………………………… TURK MAARIF COLLEGE Academic Year 2020-2021 Number:…………………………………………… 10th Forms-1st Term Chemistry Section 2 ATOM AND PERIODIC SYSTEM ATOMIC MODELS The smallest particle found in matter is called atom. Since ancient times, many theories and atomic models have been developed describe atoms. 1)Dalton Atomic Model Dalton discovered that atoms were fullfilled spherical solid particles that cannot be divided, created or destroyed. He proposed his theory by showing that there are same type of atoms in an element and different elements can combine together forming compounds. At this time-zoneprotons were discovered by Rutherford and electrons were discovered by J.J.Thomson. 2)Thomson Atomic Model (Plum Pudding Model) Thomson developed a ‘Plum Pudding Model’ and believed that atoms contain high mass of positively charged particles (protons) and small mass of negatively charged particles (electrons) which were randomly found in the positive part just like plums in a pudding. He also believed that atoms were neutral so number of positively charged particles was always equal to the number of negatively charged particles. 1 3)Rutherford Atomic Model (Nuclear Model) Rutherford investigated α-Beams Experiment (Gold Leaf Experiment) and found that there’s a positively charged part in the centre of the spherical atom and called it nucleus. He believed that protons were in the nucleus and electrons were orbiting around the nucleus. At this time-zone neutrons were discovered by Chadwick. 4)Bohr Atomic Model (Planetary Model) Bohr believed that atoms were spherical, protons and neutrons were in the nucleus of atoms and electrons were found in the circular energy levels around the nucleus at 2-D positions-just like planets around the sun. 5)Modern Theory (Quantum Model) Schrödinger and Heisenberg discovered that protons and neutrons are in the nucleus of atom while electrons are found in the specific energy levels in a cloud at 3-D positions around the nucleus. Orbitals are regions in which electrons have high probabilities to be found. 2 SUB-ATOMIC PARTICLES There are 3 types of sub-atomic particles inside an atom: protons, neutrons and electrons First sub-atomic particle discovered was electron. Electrons were discovered by J.J.Thomson. Secondly, protons were discovered by Rutherford and they were told the name ‘protons’ by Stoney. Presence of neutrons were predicted by Rutherford but discovered by Chadwick. Table 1. The charge, mass and locations of the particles in an atom. Particle Charge Mass Location Proton ____ 1 Nucleus Neutron 0 ____ Nucleus Electron -1 0 (1⁄1840 - almost zero!) In the _______ around the _______ Be CaReFuL! Remember the charge on each particle like this: p+ Protons are positive Neutrons are neutral (so electrons must be negative ones) n e- Atomic Number (Proton Number) Each element has its own atomic number which is also called proton number. It is just like ID Card number and special for each element. 3 It tells us how many protons there are in one atom of that element. Thats why it’s also called nuclear charge. Atoms of the same element have same number of protons. Different elements have different number of protons. Elements are arranged according to their atomic numbers in the Periodic Table. Atomic number = number of protons If atoms are neutral, the positive charges are cancelled out by an equal number of negative charges. So, number of protons (+) must always equal the number of electrons (-) In neutral atoms, Atomic number = number of protons = number of electrons Mass Number Mass number is sum of protons and neutrons in the nucleus of an atom. Mass number = Number of protons + Number of neutrons So, Number of neutrons = Mass number - Number of protons Example: Example: Lithium’s atomic number is 3 and its mass number is 7. How many protons, neutrons and electrons are in a lithium atom? Li Number of protons = Atomic number = 3 Number of electrons = Number of protons = 3 Number of neutrons = Mass number – Number of protons = 7 – 3 = 4 4 Atomic Symbols We can write the symbol 7 3𝐿𝑖 to show the mass number and atomic number of lithium. Atomic number is written at the bottom and to the left of the symbol. Mass number is written at the top and to the left of the symbol. Mass number Li 7 Atomic number 3 Symbol Number of neutrons = 7-3 = 4 Exercise: Find number of protons, neutrons and electrons for Atomic Symbol Number of protons Number of neutrons 40 20𝐶𝑎 and 31 15𝑃 . Number of electrons 𝟒𝟎 𝟐𝟎𝐂𝐚 𝟑𝟏 𝟏𝟓𝐏 Positively and negatively charged particles are called ions. The number written at upright of the element symbol is called ‘Net Charge’ Ions can be monatomic such as S2- and Mg2+ and polyatomic such as NH4+ and CO32-. 5 When an atom loses electron, it gets positive charge and becomes a cation. Ex.: 13Al3+ has 13 protons and 13-3=10 electrons. When an atom gains electron, it gets negative charge and becomes an anion. Ex.:8O2- has 8 protons and 8+2=10 electrons. Neutral atoms, have same number of protons and electrons. Cations have more protons than electrons since they lose electrons Anions have more electrons than protons since they gain electrons. Exercise: 56 X2+ has 32 electrons. How many neutrons are there in its 24 X2- has 24 neutrons. How many electrons does it have? nucleus? Exercise: Exercise: 6 Isotopes Atoms of an element with same number of protons and different number of neutrons are called isotopes. Since isotopes have same number of protons, they have same atomic numbers. Isotopes have different number of neutrons. That’s why they have different mass numbers. Since isotopes have different masses, they have different physical properties. For example, the ones with larger mass will have higher melting and boiling points. Isotopes have same number of electrons, so they have similar chemical properties since they are in same group of the periodic table and react similarly to each other. Exercise: Find number of protons, neutrons and electrons in these two isotopes of chlorine: 35 17𝐶𝑙 and 37 17𝐶𝑙 Atomic Symbol Number of protons Number of neutrons Number of electrons 𝟑𝟓 𝟏𝟕𝐂𝐥 𝟑𝟕 𝟏𝟕𝐂𝐥 Isotones: Elements that have same number of neutrons but different number of protons are called isotones. Example: 20 12𝑋 and 17 9𝑌 They both have 8 neutrons while their atomic and so proton numbers are different! 7 Isobars: Elements that have mass number but different atomic number are called isobars. Example: 𝟐𝟐 𝟏𝟑𝒁 and 𝟐𝟐 𝟒𝑻 They both have 22 as mass number while their atomic numbers are different! Isoelectronics: Different elements that have same number of electrons are called isoelectronics. 𝟏𝟔 2- 𝟐𝟎 𝟖𝑶 , 𝟏𝟎𝑵𝒆 and 𝟐𝟒 2+ 𝟏𝟐𝑴𝒈 They both have 10 electrons while their atomic numbers are different so they are different elements! They just have same electronic configurations since some of them are ions! Exercise.Look at the groups of elements below and decide if they are isotopes, isotones, isobars, or isoelectronics. a) 35 17𝐶𝑙 and b) 24 10𝑋 and 37 17𝐶𝑙 24 15𝑌 c) 168𝑂2− and 39 3+ d) 13 𝐴𝑙 , : _______________________ : _______________________ 24 2+ 12𝑀𝑔 14 3− 7𝑁 and : _______________________ 23 + 11𝑁𝑎 : _______________________ PERIODIC SYSTEM Periodic Table Periodic Table is a list of all of the elements. It’s a little like the alphabet of chemistry. The periodic table tells us several things… Elements were first arranged according to their mass numbers by Mendelev and Meyer.. But this was not the right one! The right version was developed by Moseley in which elements were arranged according to their atomic numbers. Appearance of a Oxygen atom in the Periodic Table is shown in the diagram below. Mass number (number of protons + number of neutrons) Atomic number (number of protons) O 16 8 8 Atomic Number: 8 Mass number: 16 Number of protons: 8 Number of electrons: 8 Number of neutrons: 16-8=8 The Arrangement of the Electrons Electrons move in the shells or energy levels around the nucleus. Each energy level can hold a certain number of electrons. Inner energy levels are always filled before outer ones. 1st shell can hold 2electrons. 2nd shell can hold 8 electrons, 3rd shell can hold 8 electrons and but it can expand upto 18 electrons. 4th shell can hold 8 electrons. 3rd shell can hold 8 electrons but it can expan upto 18 electrons 2nd shell can hold 8 electrons 1st shell can hold 2 electrons The First 20 Elements in the Periodic Table 9 There are Two Ways to Represent Atomic Structure: A. Electronic Configurations With electronic configuration elements are represented numerically by the number of electrons in their shells and number of shells. So Let’s try it... How to draw a Lithium atom 1. First, look at the Periodic Table 2. Then, determine the number of protons and hence the number of electrons. 7 (There are _3_ protons, and so _3_ electrons in a neutral lithium atom) 3. Arrange the electrons in levels, always filling up an inner (lower energy) level. 3 Li These will be arranged _2_ in the first level and _1 in the second level. This is written as _2 , 1_. Write the electronic configuration for the following elements; 𝟒𝟎 a) 𝟐𝟎𝐂𝐚=2,8,8,2 b) d) 𝟐𝟑 𝟏𝟏𝐍𝐚=2,8,1 𝟏𝟔 𝟖𝐎=2,6 e) 𝟑𝟓 𝟏𝟕𝐂𝐥=2,8,7 𝟐𝟖 c) 𝟏𝟒𝐒𝐢=2,8,4 f) 𝟏𝟏𝟓𝐁=2,3 10 B) Dot & Cross Diagrams With dot & cross diagrams elements and compounds are represented by dotsor crosses to show electrons, and circles to show the shells. For example; 14 N 7 O 16 8 35 17 Cl 8 2,6 2.8.7 11 Draw diagrams to show the arrangement of the electrons in 𝟏𝟐 a) Carbon atom - 𝟔𝐂 =2,4 𝟑𝟎 b) Phosphorus atom- 𝟏𝟓𝐏 𝟐𝟎 c) Neon atom - 𝟏𝟎𝐍𝐞 12 In the Periodic Table, vertical columns are called __groups___ and horizontal rows are called _period__ Elements were first divided into groups of 3 members-according to their similarities as called ‘Triats Rule’ by Döbereiner. Then, they were divided into groups of 8 as called ‘Octaves Rule’ by Newlands. In modern periodic table, elements are arranged in 18 groups and 7 periods, Groups contain elements with ___similar chemical__ properties. Elements in the same group have the same number of __electrons___ in the outer levels. The number of ____electrons___ in the outer level is same as the __group__ number. Group 1 6th Period Barium is in Group 2 so it has _2_ electrons in its outer level. The group 0 elements are known as the __noble__ __gases__ because they are almost completely ___stable___. thought of as being ‘full’ levels. Number of energy levels (shells) of an atom is the _period_ of it. Calcium is in 4th period so it has _4_ energy levels (shells) around its nucleus. Group 8 4th Period 13 Classification of Elements According to Their Place In Periodic Table Left part of the periodic table is mostly metals while right part is non-metals. Group 1,2, and most of group 3 elements are metals while group 5,6 and 7 elements are mostly non-metal. There are some elements that show both metallic and nonmetallic properties as well-they are called semi-metals or metalloids. (B, Si, Ge, As, At, Te, Po, Sb) Noble gases (Group 8 or 0) don’t show any metallic or nonmetallic property since they are very unreactive, have full outer shell and monatomic gases at room temperature. APPEREANCE MALLEABILITY MELTING AND BOILING POINT CONDUCTIVITY STATE AT RTP CHARGE METALS Shiny Malleable(Easy to change shape) High NONMETALS Dull (Matte) Brittle Good conductor of heat and electricity Solid except mercury which is liquid Gets positive charges Poor conductor of heat and electricity Solid, liquid or gas 14 Low Gets negative charges DUCTILITY HARDNESS COMPOUNDS OXIDES Ductile(Easy to pull into wire) Hard Forms ionic compounds with nonmetals Basic oxides Brittle Usually soft except carbon in structure of diamond Forms ionic compounds with metals and covalent compounds with nonmetals Acidic oxides GROUP 1-Alkali Metals They have 1 electron in their outer shell, want to lose 1 electron to have a full outer shell so get 1+ charge. They are all metals at room temperature. GROUP 2-Alkaline Earth Metals They have 2 electrons in their outer shell, want to lose 2 electron to have a full outer shell so get 2+ charge. They are all metals at room temperature. TRANSITION METALS Between group 2 and 3, there are transition metals that form colourful compounds and can get different electrical charges such as iron, copper, nickel, gold, silver, zinc, etc.. GROUP 7-HALOGENS They have 7 electrons in their outer shell, want to gain 1 electron to have a full outer shell so get 1- charge. They are all diatomic non-metals at room temperature that can be solid, liquid or gas. GROUP 8-NOBLE GASES They have full outer shells with 2 or 8 electrons so don’t want to lose or gain any electrons, they are stable, very unreactive monatomic gases at room trmperature. HeliumUsed in balloons NeonUsed in advertising lights ArgonUsed in light bulbs XenonUsed in lasers 15 TRENDS IN THE CHANGE OF PERIODIC PROPERTIES There’s a regular change in some physical and chemical properties of elements within a group or period. Some of these properties are; 1. Metallic and nonmetallic properties 2. Atomic volume or size 3. Ionisation energies 4. Electron affinity 5. Electronegativity 1.Metallic and nonmetallic properties: When you go down a group, metallic properties increase since metals get stronger as number. of protons and their nuclear attraction power increase. While, when you go across a period metallic properties decrease since there are more non-metals and finally noble gases! As metallic properties increase, non-metallic properties decrease! That’s why nonmetallic properties decrease down a group and increase across a period. Example. Arrange metallic properties of following elements in ascending order: 12Mg, 13Al, 19K, 20Ca Example. Arrange nonmetallic properties of following elements in descending order: 3Li, 8O, 17Cl 2.Atomic volume or size Atomic size is distance between nucleus and outermost shell of an atom. Atomic size increases as you go down a group since no. of shells increase. When you go right a period, atomic size decrease since no. of shells stay same but no. of protons and attraction force of nucleus increase and it gets outer electron more. 16 When we look at isoelectronics, the greater the nuclear charge, the smaller ionic radius since as nuclear charge attracts atoms get smaller! 😊 Example. Arrange atomic sizes of following elements in ascending order: 9F, 5B, 19K, 13Al Example. Arrange nonmetallic properties of following ions in descending order: 10Ne, 8O 2- , 7N3, 13Al3+, 2+12Mg 3.Ionisation energies 1st ionisation energy is amount of energy to remove 1 electron from gaseous state of an element. To remove 2nd electron, we need 2nd ionisation energy and so on.. X(g) + I1 X+(g) + e- X+(g) + I2 X2+(g) + e- X2+(g) + I3 X3+(g) + e- As you go down a group, ionisation energy always decrease. Because, as atomic radius gets larger, its gets easier to remove an electron from outer shell since nuclear attraction between nucleus and outer electron decrease. So, we need less energy to overcome these forces. 17 As you go across a period, ionisation energy usually increase. Because, as atomic radius gets smaller it gets harder to remove an electron from outer shell since nuclear attraction increase. So, we need more energy to overcome these forces. But; there are 2 exceptions between groups 2-3 and 5-6. Because, between these groups electrons also change their orbitals and we need less energy to remove them suddenly. So the correct order is this for ionisation energy 1A<3A<2A<4A<6A<5A<7A<8A Example. Arrange ionisation energies of following elements in ascending order: 11Na, 15P, 16S, 17Cl Example. Arrange ionisation energies of following elements in descending order: 17Cl, 16S, 13A, 12Mg Be CaReFuL! 😊 You can understand group of a substance by looking at its different ionisation energy values. If there’s a big jump between 2 ionisation energies, that means at that point the electron changes its energy level and now its not at the outer level more, it’s closer to the nucleus, it requires more energy to remove the electron so the number of ionisation energy at first step of jump is group number. 18 Example.Look at the ionization energies of element X shown below and decide its group number. Explain your answer. Ionisation energies (kJ/mol) 1st IE 2nd IE 3rd IE 4th IE 5th IE Element X 1260 2260 6940 7300 8500 ____________________________________________________________ ____________________________________________________________ ____________________________________________________________ ____________________________________________________________ 4.Electron affinity It is the amount of energy given out when an electron is added to the gaseous state of an element. X(g) + e- X-(g) + E It is similar to ionisation energy and decreases down a group, increases acrossa period. In a period, halogens have the highest electron affinity since noble gases don’t have any electron affinity value. Example. Arrange electron affinities of following elements in increasing order: 12Mg, 14Si, 15K, 18Ar 19 5.Electronegativity A measure of the tendency of an atom to attract electrons in a molecule is called electronegativity. It’s influenced by same factors with electron affinities. In general, the greater electron affinity of an atom; the greater electronegativity of an atom. F, O, N are elements with strongly high electronegativities. On the other hand, inert gases have no or little electronegativities. Example. Arrange electronegativities of following elements in decreasing order: 5Ne, 8O, 7N, 13Al, 15P SUMMARY Properties that increase down a group and decrease across a period Atomic size, metallic properties Properties that decrease down a group and increase across a period Ionisation energy, nonmetallic properties, electron affinity, electronegativity (There are exceptions) GOOD LUCK! 😊 20