CAPE CHEMISTRY MODULE 1

HOW THEORIES ARE ACCEPTED-

According to the CAPE syllabus, theories are accepted through a systematic process that involves the following stages:

1. Observation: Gathering data and making observations to identify patterns or phenomena that require explanation.

2. Hypothesis Formation: Formulating a hypothesis to provide a preliminary explanation of the observed phenomena.

3. Testing: Conducting experiments or further observations to test the hypothesis against empirical evidence.

4. Peer Review: Submitting findings to the scientific community for scrutiny and validation.

5. Reproducibility: Ensuring that results can be consistently reproduced by other researchers under similar conditions.

6. Acceptance: If a theory withstands rigorous testing and peer review, it may be accepted as a valid explanation of the phenomena until new evidence suggests otherwise.

Daltons theory

Dalton's Theory: Proposed by John Dalton in the early 19th century, this theory laid the groundwork for modern chemistry. Key tenets include:

1. All matter is made up of atoms: Atoms are indivisible particles that cannot be created or destroyed.

2. Atoms of a given element are identical: Atoms of the same element share the same properties, such as mass and size.

3. Compounds are formed by combining atoms: Atoms of different elements can combine in fixed ratios to form compounds.

4. Chemical reactions involve rearrangement of atoms: During a chemical reaction, atoms are rearranged to form new substances, but the total number

5. of atoms remains constant.

   Description of atomic structure

Atomic Structure: The atomic structure consists of a nucleus, which contains protons and neutrons, surrounded by electrons that orbit the nucleus in various energy levels.

• Nucleus: Central part of the atom, consisting of:

  • Protons: Positively charged particles that determine the atomic number of an element.mass-1 +1

  • Neutrons: Neutral particles that contribute to the atomic mass but do not affect the charge. 0 mass-1

• Electrons: Negatively charged particles that occupy electron shells or energy levels around the nucleus. mass- 1/1836 -1

  • Energy Levels: Electrons are arranged in specific layers, with each level holding a certain maximum number of electrons.

  • Electrons can move between energy levels by absorbing or releasing energy.

Overall Structure: Atoms are electrically neutral when the number of protons equals the number of electrons. This arrangement allows for chemical reactions, where atoms can bond and form compounds.

Radioactivity is a phenomenon where unstable atomic nuclei lose energy by emitting radiation. This process occurs naturally in some elements, such as uranium and thorium. As these unstable nuclei decay, they release particles or electromagnetic waves. The types of radiation released can include alpha particles, beta particles, and gamma rays.

The process of radioactivity occurs over varying time scales depending on the substance and is characterized by its half-life, the time required for half of the radioactive atoms in a sample to decay. Radioactive decay can result in the transformation of one element into another, leading to a change in the composition of the material involved.

Radioactivity is harnessed in various applications, including medical imaging and treatments, power generation in nuclear reactors, and even in dating ancient artifacts through techniques like carbon dating.

equation for decay-

alpha 223 88X=219 86 RN + 4 2 He

beta 14 6C=14 7 N + 0 -1E

gamma decay 37 18 Ar + 0 -1E= 37 17Cl

relative atomic mass of an element given isotopic masses and abundances

• Mass Number (A): The mass number of an atom is the total number of protons and neutrons present in its nucleus. It is an important indicator of the isotope of an element and is usually denoted as A.

• Isotopes: Isotopes are variants of a particular chemical element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers. For example, carbon-12 and carbon-14 are isotopes of carbon, with 6 and 8 neutrons respectively.

• Relative Atomic Mass: The relative atomic mass of an element is a weighted average of the masses of all the isotopes of that element compared to one-twelfth of the mass of a carbon-12 atom. This value takes into account the abundance of each isotope in nature.

• Isotopic Masses: Isotopic masses refer to the mass of a specific isotope of an element, usually expressed in atomic mass units (amu). Isotopic masses are measured more precisely and are crucial for calculations involving relative atomic mass.

calculation of relative atomic mass of an element ,given isotopic masses and abundances.

To determine relative atomic mass, we simply multiply each isotopic mass by its abundance, add all the values together and divide the total value by 100 percent.

explain how data from emission spectra provide evidence for discrete energy levels within the atom-

When an atom gets energy, its tiny particles called electrons can jump to higher places, like a kid jumping up to a higher step on a playground. When the electrons get tired, they come back down to where they were before and give off some energy as light. This light is made up of different colors, and instead of being a smooth rainbow, we see specific colors, like red or blue, appearing as lines. These lines show that the electrons have only certain spots they can jump to, making it look like there are steps instead of just a big slide. So, this means that atoms have specific energy levels, kind of like a ladder.

The electrons in atoms have certain fixed values of energy. The smallest fixed amount of energy required for a change is called a quantum of energy. The energy is discrete or quantised, like the energy levels which the electrons are arranged in. The atom is most stable when electrons are in the lowest energy levels possible for each of them.

when they become excited, they move to higher levels and fall back to lower emitting radiation of characteristic frequency. This is the origin of the line of emission spectrum. The energy difference between any two energy levels is given by totalE=hv planck’s constant - 6.63× 10^-34js^-1

Atomic emission spectra -when electrical or thermal energy is passed through a gaseous sample of an element, the radiation is emitted only at certain wave lengths or frequencies. An emission spectrum differs from a normal visible light spectrum in that it is made up of separate lines, the lines converge as their frequency increases.

each line in the hydrogen emission spectrum is a result of electrons moving from a higher to a lower energy level. Among several series of lines seen are the -lyman series- previously excited electrons fall back to the n=1 energy level(seen in ultaviolet region)

balmer series seen in the visible region wher previously excited electrons fall back to the n=2 energy level.

IONISATION ENERGY.

Factors that Influence First Ionization Energy of Elements:

1. Atomic Radii: As the atomic radius increases, the distance between the nucleus and the outermost electron increases, resulting in a weaker attraction. This leads to lower ionization energy.

2. Nuclear Charge: The number of protons in the nucleus determines the nuclear charge. A higher nuclear charge creates a stronger attraction between the nucleus and the electrons, increasing ionization energy.

3. Shielding Effect: Inner-shell electrons shield outer-shell electrons from the full effect of the nuclear charge. Increased shielding reduces the effective nuclear charge experienced by outer electrons, leading to lower ionization energy.

4. ionization energy- energy needed to remove one electron from each atom n one mole of atoms of an element in its gaseous state to form one mole of gaseous ions.

ionization energy data for elements in Period 3 provide evidence for the presence of subshells through observable trends in ionization energy values. As we move across Period 3 from sodium (Na) to argon (Ar), the ionization energy generally increases due to the greater nuclear charge, which enhances the attraction between the nucleus and the outer electrons. However, there are notable exceptions that highlight the effect of subshells:

1. Sodium to Magnesium (Na to Mg): As we move from Na to Mg, the ionization energy increases because the outer electron in magnesium is in the 3s subshell, where there is less shielding.

2. Magnesium to Aluminum (Mg to Al): There is a drop in ionization energy when moving from Mg to Al. This decrease can be attributed to the filling of the 3p subshell in aluminum. The additional electron in the 3p subshell experiences increased shielding from the inner electrons and is further away from the nucleus than the electron in the 3s subshell, making it easier to remove.

3. Phosphorus to Sulfur (P to S): A similar trend occurs between phosphorus and sulfur, where there is again a drop in ionization energy due to electron-electron repulsion when two electrons occupy the same 3p orbital in sulfur. This repulsion reduces the energy needed to remove an electron, confirming the subshell structure.

Overall, these variations in ionization energy illustrate the existence of subshells and the impact of electron configuration on the stability and reactivity of elements within Period 3.

1. Ionic Bonds: Attractive forces between positively and negatively charged ions. These bonds form when electrons are transferred from one atom to another, resulting in the formation of ions.

2. Covalent Bonds: Bonds formed when two atoms share one or more pairs of electrons. This sharing allows both atoms to achieve a stable electronic configuration.

3. Hydrogen Bonds: A type of dipole-dipole attraction that occurs between molecules when hydrogen is bonded to a strongly electronegative atom (such as oxygen, nitrogen, or fluorine). This creates a partial positive charge on the hydrogen atom, which can then attract other electronegative atoms in nearby molecules.

4. Metallic Bonds: The forces of attraction between metal atoms and the surrounding sea of delocalized electrons. This bonding gives metals their characteristic properties, such as conductivity and malleability.

5. Van der Waals Forces: Weak attractions between molecules or parts of molecules that result from transient local partial charges. These forces include dipole-dipole interactions and London dispersion forces, and they play a significant role in the physical properties of substances.

The relationship between forces of attraction and states of matter can be summarized as follows:

1. Solids: In solids, particles are closely packed together and experience strong forces of attraction. These forces keep the particles in fixed positions, resulting in a definite shape and volume. The strong intermolecular forces lead to low kinetic energy among particles, causing them to vibrate in place rather than move freely.

2. Liquids: In liquids, particles are still in close contact but have more kinetic energy than in solids. The forces of attraction between particles are moderate, allowing them to slide past one another. This gives liquids a definite volume but no definite shape, as they take the shape of their container.

3. Gases: In gases, particles are far apart and experience weak forces of attraction. The kinetic energy of gas particles is much higher than in solids and liquids, allowing them to move freely and occupy the entire volume of their container. This results in no definite shape or volume.

Overall, as the strength of the forces of attraction decreases from solids to liquids to gases, the kinetic energy of the particles increases, leading to changes in the physical state of matter.

The physical properties of matter, such as melting points, boiling points, and solubilities, can be attributed to the differences in the strength of forces of attraction between particles:

1. Melting Points:

   • Strong Forces (e.g., Ionic and Covalent Bonds): Substances with strong forces of attraction, such as ionic compounds and covalent network solids, typically have high melting points because a significant amount of energy is required to overcome these strong attractions.

   • Weak Forces (e.g., Van der Waals Forces): Substances with weaker forces of attraction, such as those bonded by van der Waals forces, usually have lower melting points, as less energy is needed to break these bonds.

2. Boiling Points:

   • Ionic and Hydrogen Bonds: Liquids with strong forces, like ionic compounds or those with hydrogen bonds, have higher boiling points because it takes more energy to separate the molecules and transition into a gaseous state.

   • Covalent Compounds and Van der Waals Forces: Substances that are held together by weaker van der Waals forces or covalent bonds will have lower boiling points since less energy is needed for the transition to the gas phase.

3. Solubilities:

   • Like Dissolves Like: The ability of a substance to dissolve in a solvent is largely influenced by the nature and strength of the forces of attraction between solute and solvent molecules. For instance, polar solvents (like water) tend to dissolve ionic and polar covalent substances due to strong dipole-dipole interactions or ion-dipole interactions.

   • Nonpolar Substances: Nonpolar substances tend to be soluble in nonpolar solvents as both have similar weak van der Waals forces. Thus, substances with stronger attractions are less likely to dissolve in solvents with weaker attractions.

Overall, the variation in melting points, boiling points, and solubilities of substances are directly related to the strength of the intermolecular or ionic forces of attraction present in the material.

Formation of Bonds:

(i) Ionic Bonds:

Ionic bonds form when electrons are transferred from one atom to another, leading to the creation of ions. Typically, this occurs between metals and nonmetals.

• Metal Atoms: Metals tend to lose electrons, becoming positively charged cations.

• Non-Metal Atoms: Nonmetals gain these electrons to become negatively charged anions.

• Attraction: The electrostatic attraction between the positively charged cations and negatively charged anions results in the formation of ionic bonds. The resulting structures are ionic compounds, often characterized by high melting and boiling points.

(ii) Covalent Bonds:

Covalent bonds form when two atoms share one or more pairs of electrons to achieve a stable electronic configuration.

• Orbital Overlap: This involves the overlap of atomic orbitals from each atom.

  • Sigma (σ) Bonds: A σ bond forms when the orbitals overlap end-to-end. This bond allows for free rotation around the bond axis and is the strongest type of covalent bond.

  • Pi (π) Bonds: A π bond forms when the overlap occurs sideways, typically involving p orbitals. π bonds are generally weaker than σ bonds and allow for restricted rotation due to their spatial orientation. Therefore, in double or triple bonds, one bond is always a σ bond while the additional bonds are π bonds.

(iii) Metallic Bonds:

Metallic bonds occur in metals where atoms are packed closely together in a lattice structure.

• Lattice of Positive Ions: In this structure, metal atoms lose their valence electrons, becoming positively charged ions.

• Mobile Electrons: The lost electrons are not associated with any specific atom and move freely throughout the lattice, forming a 'sea of delocalized electrons.' This mobility allows metals to conduct electricity and heat efficiently.

• Properties: Metallic bonds give metals their characteristic properties, such as malleability, ductility, and luster.

Electronegativity and Polarity of Bonds:

• Electronegativity: This refers to the tendency of an atom to attract electrons in a chemical bond. Differences in electronegativity between bonded atoms can influence bond type:

  • Ionic Bonds: Form between atoms with a large difference in electronegativity (typically greater than 1.7).

  • Polar Covalent Bonds: Involves a moderate difference in electronegativity, resulting in an unequal sharing of electrons and the formation of partial charges (δ+ and δ-).

  • Nonpolar Covalent Bonds: Occur when two atoms of the same element share electrons equally (no difference in electronegativity).

Understanding these concepts helps clarify how different types of bonds form and influence the chemical and physical properties of substances

Melting and Boiling Points

• Ionic Compounds:

  • High Melting/Boiling Points: Ionic compounds typically have high melting and boiling points due to the strong electrostatic forces of attraction between the oppositely charged ions in the crystal lattice.

  • Example: Sodium chloride (NaCl) has a melting point of 801 °C and a boiling point of 1413 °C.

• Covalent Compounds:

  • Lower Melting/Boiling Points: Covalent compounds usually have lower melting and boiling points compared to ionic compounds because the intermolecular forces (like van der Waals forces) are generally weaker than ionic bonds.

  • Example: Water (H₂O) has a melting point of 0 °C and a boiling point of 100 °C.

Solubility in Polar and Nonpolar Solvents

• Ionic Compounds:

  • Soluble in Polar Solvents: Ionic compounds are generally soluble in polar solvents (like water) because the polar molecules can stabilize the ions when they dissociate.

  • Example: NaCl readily dissolves in water due to ion-dipole interactions.

• Covalent Compounds:

  • Solubility Depends: Covalent compounds can vary in solubility depending on their polarity. Polar covalent compounds can dissolve in polar solvents, while nonpolar covalent compounds dissolve in nonpolar solvents.

  • Example: Sugar (C₁₂H₂₂O₁₁) is soluble in water (polar), while oil (nonpolar) does not dissolve in water.

Electrical Conductivity

• Ionic Compounds:

  • Conductive When Dissolved or Molten: Ionic compounds conduct electricity when they are dissolved in water or melted because the ions are free to move and carry charge.

  • Example: A NaCl solution conducts electricity when an electric current is applied.

• Covalent Compounds:

  • Usually Non-Conductive: Covalent compounds generally do not conduct electricity in either solid or dissolved form because they do not contain charged particles (ions) that can move to carry an electrical current.

  • Example: Sugar solution does not conduct electricity; it lacks free-moving ions.

the origin of inter-molecular forces- intermolecular forces arise because of the attraction between dipoles in neighbouring molecules .There are three types of intermolecular forces.

permanent dipole-dipole forces - weak attractive forces between partial positive of the dipole of one molecule and the partial negative of the dipole of the neighbouring molecule.

van der waals forces - not permanent all atoms and molecules , including noble gas atoms have van der waals forces.

hydrogen bonging - special form of permanent dipole bonding requires h bonded to f oor n tom. these are most electronegative atoms. a second molecule having a f o n atom with a lone pair of electrons.