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98 Cards in this Set
- Front
- Back
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Define Atom |
Smallest particle of matter |
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Define Vallency |
Amount of electrons that can be gained or lost by an atom during a reaction |
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Define Isotopes |
Atoms of the same element with same no. of protons but different no. of nutreons |
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Why do Isotopes have same chemical properties? |
They have same no. of electrons in the outermost shell |
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Mass of electron |
1/1836 amu |
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What does nuclear charge mean? |
No of protons in nucleus |
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Shape of S subshell |
Spherical |
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S subshell's nature |
Non directional in nature |
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Shape of P subshell |
Dumbbell |
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Order of increasing energy of electrons (1S2 representation) |
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s |
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We name shells n=1, n=2 What does n mean? |
Principle quantum number |
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Why are group 1 and group 2 called s-block? |
Because valence electron enters into s-subshell |
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General formula of s-block |
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General formula of p-block |
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What happens to atomic size across a period and why? |
Decreases •nuclear charge increases •electrons added to the same shell, so nuclear force of attraction on electrons increases |
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What happens to atomic size down the group and why? |
Increases •Electrons are added to the new shell so shielding effect increases •Nuclear force of attraction on the electrons decreases |
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Why are cations smaller than neutral atom? |
Same no. of protons have to attract less no. of electrons so effective nuclear force of attraction on electron increases, making ion smaller. |
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Why are size of anions larger than neutral atom? |
Newly added electrons are repelled by valence electrons which causes expansion of the ion. |
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Define Ionization energy |
Amount of energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form gaseous unipositive ions |
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Ionization process is always: Endothermic or exothermic |
Endothermic |
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What happens to IE across the period? And Why? |
Increases •Nuclear charge increases •Atomic size decreases •Shielding effect remains the same as electrons are added to the same effect so nuclear force of attractin on the outer electrons will be more so IE increases |
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What happens IE down the group? And why? |
Decreases •Atomic size increases •Shielding effect increases as electrons are added to new shells •Increase in nuclear charge is cancelled by the increase in shielding effect, so nuclear force of attraction on outer electrons decrease |
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Be to B IE1 decreases. Explain |
In Boron, the valence electron is in 2p atomic orbital which is at a higher energy level away from the nucleus and experiences more shielding so nuclear force of attraction will be less |
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N to O IE decreases. Explain |
In N all the 2p orbitals are half-filled whereas in O, one of the 2p orbitals have paired electrons and they repel. This is called spin pair repulsion. This makes removal of electron easy |
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Define Second Ionization energy |
Amount of energy required to remove 1 electron from each ion of one mole of gasueos unipositive ions to give one mole of gaseous dipositive ions |
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Define successive ionization energy |
The total energy required to remove the electrons one by one mole of gaseous atom is called successive IE |
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Define Relative atomic mass, Ar |
It is the weighted average mass of all atoms of an element(or all the isotopes of an element) compared to 1/12th the mass of a Carbon-12 atom |
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Relative molecular mass, Mr |
It is the weighted average mass of atoms of a molecule compared to 1/12th the mass of a Carbon-12 atom |
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Relative formula mass |
It is the weighted average mass of one formula unit of a molecule compared to 1/12th the mass of a Carbon-12 atom |
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Relative Isotopic mass |
It is the mass of an isotope of an element compared to 1/12th of a Carbon-12 atom |
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Define Mole |
It is the amount of substance that contain same no.of particles as in 12g of Carbon-12(6.02x10^23 atoms) |
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Define Empirical formula |
It is the simplest whole no. ratio of atoms of each element present in a molecule |
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Define Molecular formula |
It is the actual no. of atoms of each element present in a molecule |
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Define Mass Spectrum |
It is the graphical representation of amount of each isotope relative to its percentage abundance |
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Combustion of hydrocarbon general formula |
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Define Ionic bonding |
The electrostatic force of attraction between positive and negative ions |
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Define lattice |
3 dimensional regular arrangement of particles |
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State the properties of Ionic Compounds |
•Soluble in polar solvents and insoluble in organic solvents •Insulators in solid state •Conductors in aqueous and molten form •Brittle •High melting point |
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Define Coordinate number |
The total no. of oppositely charged ions present around the central ion in an ionic lattice |
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Define covalent bonding |
The electrostatic force of attraction between the nuclei of the bonded atoms and the shared pair of electrons between them |
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Define electron deficient compounds |
Compounds in which the central atom is not attaining octet configuration even after bonding |
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Define Electron rich compounds/ Expanded octet |
The compounds in which the central atom have more than 8 electrons after bonding |
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What are conditions of coordinate bonding? |
1. At least one lone pair in donor atom 2. The acceptor should have incomplete octet |
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Temp. At which AlCl3 combines to form Al2Cl6 |
180°C to 190°C |
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Define bond angle |
The angle between orbitals containing bond pair of electrons |
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Define isoelectronic species |
The atoms or ions with same no. of electrons |
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Na to Al, melting point increases. Explain |
Na to Al, No. of delocalised ē increases in the metal lattice and charge density of metal ion increases so, the electrostatic force of attraction between metal ions and electrons become stronger which makes the metallic bond stronger so melting point increases |
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Why does Silicon have the highest melting point? |
Silicon have gaint covalent structure. It requires large amount of energy to break covalent bonds |
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Define Metallic bonding |
The metal lattice consists of positive metal ions in a sea of delocalized ē with strong electrostatic force of attraction between them |
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Factors affecting strength of metallic bond |
No. Of delocalised ē Charge density |
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Define electronegativity |
The ability of an atom in a covalent bond to attract the shared pair of ē towards itself |
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Define bond polarity |
The unequal sharing of ē pairs in a covalent bond due to the difference in Electronegativity of the bonded atoms |
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Most electronegative element |
Flouride |
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Least electronegative element |
Francium |
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Define bond energy |
Amount of energy required to break 1 mole of a particular bond in a molecule in the gasueos state |
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Define Activation energy |
Minimum amount of energy required for a reaction to take place. |
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Standard Molar Enthalpy change |
It is the enthalpy change when the reactants react to form products according to the balanced chemical equation |
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Standard Molar enthalpy change of formation |
It is the enthalpy change when 1 mole of a compound is formed from its elements in their standard state under standard conditions. |
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Standard Molar enthalpy change of combustion |
It is the enthalpy change when 1 mole of a substance to completely burnt in excess Oxygen |
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Standard Molar enthalpy change of atomization |
Enthalpy change when 1 mole of gaseous is formed from its element in its standard state under standard condition |
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Standard Molar enthalpy change of neutralization |
Enthalpy change when 1 mole of H+ ions from an acid completely neutralized by a base under standard conditions to give 1 mole of H2O |
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Standard Molar enthalpy change of solution
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It is the enthalpy change when 1 mole of a solute dissolved in a solvent to form an infinitely dilute solution under standard conditions |
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Standard Molar enthalpy change of hydration |
It is the enthalpy change when 1 mole of gaseous ions of an element dissolves in water to form hydrated ions in infinitely dilute solution |
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Bond length |
The average distance between the nuclei of two bonded atoms |
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The factors affecting bond length |
When atomic size increases, bond length increases When orbital overlap increases, bond length decreases No. of bonds between the atoms increases, bond length decreases |
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The factors affecting bond energy |
When atomic size increases, bond energy decreases When orbital overlap increases, bond energy increases No. of bonds between the atoms increases, bond energy increases |
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Define Functional group |
Atoms or groups of atoms which gives characteristic properties to an organic compound |
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Define Structural isomerism |
Compounds having same molecular formula but have different structural formula |
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Define Chain isomerism |
Have same molecular formula but differ in the length of Carbon chain |
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Define Position isomerism |
Isomers with same molecular formula but differ in the position of functional group |
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Define Functional group isomerism |
Isomers with same molecular formula but different functional group |
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Define Geometrical isomerism |
Compounds with same molecular formula and structural formula but different arrangement of atoms about C=C due to the restricted rotation |
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Define Stereoisomerism |
Compounds with same molecular formula and structural formula but different arrangement of atoms in space |
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Define Optical isomerism |
Compounds with the same molecular and structural formula but can have Non-superimposable mirror image form and at least one chiral Carbon atom |
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What is a Dextrorotatory isomer? |
Optical isomer which rotates the plane of polarized light towards the right side. |
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What is a Laevorotatory isomer?
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Optical isomer which rotates the plane of polarized light towards the left side. |
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What is Racemic mixture? |
A mixture containing 50% Dextro isomer and 50% Laevo isomer will be optically inactive. |
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What are Nucleophiles? |
Nucleophiles are molecules or ions which can donate a pair of electrons |
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What are Electrophiles? |
Molecules or ions which can accept a pair of electrons |
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Define Homolytic fission |
Fission of a covalent in which the atoms leave with its own electron so that free radicals are formed |
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Define Heterolytic fission |
Fission of a covalent bond with unequal distribution of electrons so ions are formed |
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Define Hess' law |
The total enthalpy change of a reaction is independent of the steps involved provided the initial and final conditions are the same |
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Define Reversible reactions |
Reactions taking place in forward as well as backward direction |
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Define Dynamic equilibrium |
The point at which rate of the forward reaction is equal to the rate of backward reactions in a closed system |
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Le chatelier's principle |
It states that if the conditions of a system at equilibrium is changed, the system will change the position of equilibrium to oppose the effect of that change |
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Conditions for Haber process (production of NH3) |
450C 200 atm Catalyst: Iron |
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Steps involved in Contact process |
S + O2 = SO2 2SO2 + O2 = 2SO3 (reversible) SO3 + H2SO4 = H2S2O7 H2S2O7 + H20 = 2H2SO4 |
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Conditions for manufacture of SO3 |
450C 1-2 atm V2O5 (Vanadium(V) oxide) |
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Boyle's law |
At constant temperature, the pressure of a gas is inversely proportional to the volume |
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Charle's law |
At constant pressure, volume of a gas is proportional to temperature |
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Avagadro's law |
Equal volumes of all gases at similar conditions of temperature and pressure contains equal no. of molecules |
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State the kinetic theory assumptions for ideal gases |
1.Gas consists of atoms or molecules which are in random motion 2. Pressure of a gas is caused by collision of gas particles with the walls of the container 3.The collisions are perfectly elastic (no loss or gains of kinetic energy during collisions) 4. The absolute of gas is proportional to the average kinetic energy of the particles 5. There is no intermolecular force of attraction between gas particles 6. Volume occupied by a single gas particle is negligible when comparing with the total volume of the gas container |
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Define vapour pressure |
The pressure exerted by vapours on the surface of the liquid when they are in equilibrium |
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Define high vapour pressure |
If the amount of vapours is more than the liquid then the pressure exerted by the vapours on the surface of liquid when they are in equilibrium is called high vapour pressure |
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Uses of Al |
Used in food containers since it is non-toxic Used in aircraft bodies because of its low density Used in overhead electrical cables because of low density and good electrical conductivity |
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Uses of Cu |
Electrical wires because of its good electrical conductivity Hot water pipes because of it's unreactive with water Used for making cooking pans because it's a good conductor of heat |
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Uses of waste management |
Prevent atmospheric pollution The area used for waste disposal can be saved Saves money |
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Why are Alkanes unreactive? |
They are nonpolar They have high bond energy They have high Activation energy |
