How many atoms does 32 grams of sulfur contain?

Higher concentration make a reaction faster due to faster collisions. The correct answer is "There are more collisions per second."

When the concentration of reactants is increased, it leads to a higher number of particles or molecules in the reaction mixture. As a result, there are more frequent collisions between the reactant particles per unit time.

In a chemical reaction, the reactant particles must collide with sufficient energy and proper orientation for a successful reaction to occur. By increasing the concentration, the chances of successful collisions increase because there are more particles available to collide with each other.

However, it's important to note that while a higher concentration increases the frequency of collisions, it does not necessarily guarantee that all collisions will result in a reaction. The energy of the collisions and proper orientation are also essential factors for a successful reaction.

Learn more about collisions from the link given below.

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Answer:

Approximately [tex]8.3 \times 10^{-2}\; \rm mol \cdot L^{-1}[/tex].

Explanation:

The [tex]K_a[/tex] in this question refers the dissociation equilibrium of [tex]\rm HC_6H_4ClO[/tex] as an acid:

[tex]\rm HC_6H_4ClO\, (aq) \rightleftharpoons H^{+} \, (aq) + C_6H_4ClO^{-}\, (aq)[/tex].

[tex]\displaystyle K_a\left(\mathrm{HC_6H_4ClO}\right) = \frac{\left[\mathrm{H^{+}}\right] \cdot \left[\mathrm{C_6H_4ClO^{-}}\right]}{\left[\mathrm{HC_6H_4ClO}\right]}[/tex].

However, the question also states that the solution here has a [tex]\rm pH[/tex] of [tex]11.05[/tex], which means that this solution is basic. In basic solutions at [tex]\rm 25\;^\circ C[/tex], the concentration of [tex]\rm H^{+}[/tex] ions is considerably small (typically less than [tex]10^{-7}\;\rm mol \cdot L^{-1}[/tex].) Therefore, it is likely not very appropriate to use an equilibrium involving the concentration of [tex]\rm H^{+}[/tex] ions.

Here's the workaround: note that [tex]\rm C_6H_4ClO^{-}\, (aq)[/tex] is the conjugate base of the weak acid [tex]\rm HC_6H_4ClO\, (aq)[/tex]. Therefore, when [tex]\rm C_6H_4ClO^{-}\, (aq)[/tex] dissociates in water as a base, its [tex]K_b[/tex] would be equal to [tex]\displaystyle \frac{K_w}{K_a} \approx \frac{10^{-14}}{K_a}[/tex]. ([tex]K_w[/tex] is the self-ionization constant of water. [tex]K_w \approx 10^{-14}[/tex] at [tex]\rm 25\;^\circ C[/tex].)

In other words,

[tex]\begin{aligned} & K_b\left(\mathrm{C_6H_4ClO^{-}}\right) \\ &= \frac{K_w}{K_a\left(\mathrm{HC_6H_4ClO}\right)} \\ &\approx \frac{10^{-14}}{6.6 \times 10^{-10}} \\ & \approx 1.51515 \times 10^{-5}\end{aligned}[/tex].

And that [tex]K_b[/tex] value corresponds to the equilibrium:

[tex]\rm C_6H_4ClO^{-}\, (aq) + H_2O\, (l) \rightleftharpoons HC_6H_4ClO\, (aq) + OH^{-}\, (aq)[/tex].

[tex]\displaystyle K_b\left(\mathrm{C_6H_4ClO^{-}}\right) = \frac{\left[\mathrm{HC_6H_4ClO}\right]\cdot \left[\mathrm{OH^{-}}\right]}{\left[\mathrm{C_6H_4ClO^{-}}\right]}[/tex].

The value of [tex]K_b[/tex] has already been found.  

The [tex]\rm OH^{-}[/tex] concentration of this solution can be found from its [tex]\rm pH[/tex] value:

[tex]\begin{aligned}& \left[\mathrm{OH^{-}}\right] \\ &= \frac{K_w}{\left[\mathrm{H}^{+}\right]} \\ & = \frac{K_w}{10^{-\mathrm{pH}}} \\ &\approx \frac{10^{-14}}{10^{-11.05}} \\ &\approx 1.1220 \times 10^{-3}\; \rm mol\cdot L^{-1} \end{aligned}[/tex].

To determine the concentration of [tex]\left[\mathrm{HC_6H_4ClO}\right][/tex], consider the following table:

[tex]\begin{array}{cccccc}\textbf{R} &\rm C_6H_4ClO^{-}\, (aq) & \rm + H_2O\, (l) \rightleftharpoons & \rm HC_6H_4ClO\, (aq) & + & \rm OH^{-}\, (aq) \\ \textbf{I} & (?) & \\ \textbf{C} & -x & & + x& & +x \\ \textbf{E} & (?) - x & & x & & x\end{array}[/tex]

Before hydrolysis, the concentration of both [tex]\mathrm{HC_6H_4ClO}[/tex] and [tex]\rm OH^{-}[/tex] are approximately zero. Refer to the chemical equation. The coefficient of [tex]\mathrm{HC_6H_4ClO}[/tex] and [tex]\mathrm{HC_6H_4ClO}[/tex] are the same. As a result, this equilibrium will produce [tex]\rm OH^{-}[/tex] and [tex]\mathrm{HC_6H_4ClO}[/tex] at the exact same rate. Therefore, at equilibrium, [tex]\left[\mathrm{HC_6H_4ClO}\right] \approx \left[\mathrm{OH^{-}}\right] \approx 1.1220 \times 10^{-3}\; \rm mol\cdot L^{-1}[/tex].

Calculate the equilibrium concentration of [tex]\left[\mathrm{C_6H_4ClO^{-}}\right][/tex] from [tex]K_b\left(\mathrm{C_6H_4ClO^{-}}\right)[/tex]:

[tex]\begin{aligned} & \left[\mathrm{C_6H_4ClO^{-}}\right] \\ &= \frac{\left[\mathrm{HC_6H_4ClO}\right]\cdot \left[\mathrm{OH^{-}}\right]}{K_b}\\&\approx \frac{\left(1.1220 \times 10^{-3}\right) \times \left(1.1220 \times 10^{-3}\right)}{1.51515\times 10^{-5}}\; \rm mol \cdot L^{-1} \\ &\approx 8.3 \times 10^{-2}\; \rm mol \cdot L^{-1}\end{aligned}[/tex].

Answer:

C HI = 0.0968 M

Explanation:

∴ a: HI(g)

∴ order (α) = 1

∴ rate constant (K) = 7.21/s

∴ Initial concenttration HI(g) (Cao) = 0.440 M

∴ t = 0.210 s ⇒ Ca = ?

⇒ - δCa/δt = KCa

⇒ - ∫δCa/Ca = K*∫δt

⇒ Ln(Cao/Ca) = K*t

⇒ Cao/Ca = e∧(K*t)

⇒ Cao/e∧(K*t) = Ca

⇒ Ca = (0.440)/e∧((7.21*0,21))

⇒ Ca = 0.440/4.54533

⇒ Ca = 0.0968 M

Answer:

See explanation below

Explanation:

In this reaction we have the ethyl acetoacetate which is reacting with 2 eq of sodium etoxide. The sodium etoxide is a base and it usually behaves as a nucleophyle of many reactions. Therefore, it will atract all the acidics protons in a molecule.

In the case of the ethyl acetoacetate, the protons that are in the methylene group (CH3 - CO - CH2 - COOCH2CH3) are the more acidic protons, therefore the etoxide will substract these protons instead of the protons of the methyl groups. This is because those hydrogens (in the methylene group) are between two carbonile groups, which make them more available and acidic for any reaction. As we have 2 equivalents of etoxide, means that it will substract both of the hydrogen atoms there, and then, reacts with the Br - CH2CH2 - Br and form a product of an aldolic condensation.

The mechanism of this reaction to reach X is shown in the attached picture.

Answer:

0.625 mol

27.5 g

Explanation:

Given data

We can find the moles of dinitrogen monoxide gas using the ideal gas equation.

P × V = n × R × T

n = P × V/R × T

n = 0.500 atm × 30.0 L/0.0821 atm.L/mol.K × 292.2 K

n = 0.625 mol

Then, we can find the mass corresponding to 0.625 moles.

0.625 mol × 44.01 g/mol = 27.5 g

Answer:

0.626 moles of dinitrogen monoxide gas was collected, this is 27.6 grams

Explanation:

Step 1: Data given

Temperature of Dinitrogen monoxide gas = 19.0 °C = 292 K

Volume of the flask = 30.0 L

Pressure in the flask = 0.500 atm

Step 2: Calculate moles of Dinitrogen monoxide gas

p*V = n*R*T

⇒with p = the pressure of the gas = 0.500 atm

⇒with V = the volume of the flask = 30.0 L

⇒with n = the number of moles of N2O gas = TO BE DETERMINED

⇒with R = the gas constant = 0.08206 L*atm/mol*K

⇒with T = the temperature = 292 K

n = (p*V) / (R*T)

n = (0.500 * 30.0) / (0.08206 * 292)

n = 0.626 moles

Step 3: Calculate mass N2O gas

Mass N2O gas = moles N2O * molar mass N2O

Mass N2O = 0.626 moles * 44.013 g/mol

Mass N2O = 27.55 grams ≈ 27.6 grams

0.626 moles of dinitrogen monoxide gas was collected, this is 27.6 grams

Answer:

We need 1136.4 kJ energy

Explanation:

Step 1: Data given

The mass of water = 369 grams

the initial temperature = -10°C

The finaltemperature = 125 °C

Cm (ice)=36.57 J/(mol⋅∘C)

Cm (water)=75.40 J/(mol⋅∘C)

Cm (steam)=36.04 J/(mol⋅∘C)

ΔHfus=+6010 J/mol

ΔHvap=+40670 J/mol

Step 2: :Calculate the energy needed to heat ice from -10 °C to 0°C

Q = n*C*ΔT

⇒Q = the energy needed to heat ice to 0°C

⇒with n = the moles of ice = 369 grams / 18.02 g/mol = 20.48 moles

⇒with C = the specific heat of ice = 36.57 / J/mol°C

⇒ ΔT = the change of temperature = 10°C

Q = 20.48 moles * 36.57 J/mol°C * 10°C

Q = 7489.5 J = 7.490 kJ

Step 3: calculate the energy needed to melt ice to water at 0°C

Q = n* ΔHfus

Q = 20.48 moles * 6010 J/mol

Q = 123084.8 J = 123.08 kJ

Step 4: Calculate energy needed to heat water from 0°C to 100 °C

Q = n*C*ΔT

⇒Q = the energy needed to heat water from 0°C to 100 °C

⇒with n = the moles of water = 369 grams / 18.02 g/mol = 20.48 moles

⇒with C = the specific heat of water  =75.40 J/(mol⋅∘C)

⇒ ΔT = the change of temperature = 100°C

Q = 20.48 moles * 75.40 J/mol°C* 100°C

Q = 154419.2 J = 154.419 kJ

Step 5: Calculate energy needed to vapourize water to steam at 100°C

Q = 20.48 moles * 40670 J/mol

Q = 832921.6 J = 832.922 kJ

Step 6: Calculate the energy to heat steam from 100 °C to 125 °C

⇒Q = the energy needed to heat steam from 100 °C to 125 °C

⇒with n = the moles of steam = 369 grams / 18.02 g/mol = 20.48 moles

⇒with C = the specific heat of steam  = 36.04 J/(mol⋅∘C)

⇒ ΔT = the change of temperature = 25 °C

Q = 20.48 moles * 36.04 J/mol*°C * 25°C

Q = 18452.5 J = 18.453 kJ

Step 7: Calculate total energy needed

Q = 7489.5 J + 123084.8 J + 154419.2 J +  832921.6 J + 18452.5 J

Q = 1136367.6 J

Q = 1136.4 kJ

We need 1136.4 kJ energy

The amount of energy needed is [tex]\( \\1120 \text{ kJ}} \).[/tex]

To calculate the amount of energy required to change 369 g of water from ice at[tex]\(-10^\circ \text{C}\)[/tex] to steam at [tex]\(125^\circ \text{C}\),[/tex] we need to consider the different phases of water and the energy required for each phase change and temperature change.

Let's break down the calculation step by step.

1. Energy to heat ice from [tex]\(-10^\circ \text{C}\) to \(0^\circ \text{C}\).[/tex]

  First, calculate the energy required to heat the ice initially at[tex]\(-10^\circ \text{C}\)[/tex] to its melting point [tex]( \(0^\circ \text{C}\) ).[/tex]

  Specific heat capacity of ice [tex](\(C_m\)): \(36.57 \text{ J/(mol} \cdot \text{°C)}\).[/tex]

  Molar mass of water (H₂O). [tex]\(18.015 \text{ g/mol}\).[/tex]

  Amount of ice in moles:

[tex]\[ \text{moles of ice} = \frac{369 \text{ g}}{18.015 \text{ g/mol}} \approx 20.49 \text{ mol} \][/tex]

  Energy required to heat ice.

[tex]\[ Q_1 = n \times C_m \times \Delta T = 20.49 \text{ mol} \times 36.57 \text{ J/(mol} \cdot \text{°C)} \times (0^\circ \text{C} - (-10^\circ \text{C})) = 7485.39 \text{ J} \][/tex]

[tex]\( Q_1 = 7.48539 \text{ kJ} \)[/tex]

2. Energy to melt the ice at [tex]\(0^\circ \text{C}\).[/tex]

Latent heat of fusion [tex](\(\Delta H_{\text{fus}}\)) for ice: \(6.01 \text{ kJ/mol}\)[/tex].

Energy required to melt ice:

[tex]\[ Q_2 = \text{moles of ice} \times \Delta H_{\text{fus}} = 20.49 \text{ mol} \times 6.01 \text{ kJ/mol} = 123.05 \text{ kJ} \][/tex]

[tex]\( Q_2 = 123.05 \text{ kJ} \)[/tex]

3.Energy to heat water from[tex]\(0^\circ \text{C}\) to \(100^\circ \text{C}\)[/tex].

Specific heat capacity of water[tex](\(C_m\))[/tex]: [tex]\(75.40 \text{ J/(mol} \cdot \text{°C)}\)[/tex].

Energy required to heat water:

[tex]\[ Q_3 = \text{moles of water} \times C_m \times \Delta T = 20.49 \text{ mol} \times 75.40 \text{ J/(mol} \cdot \text{°C)} \times (100^\circ \text{C} - 0^\circ \text{C}) = 155,297.80 \text{ J} \][/tex] [tex]\( Q_3 = 155.2978 \text{ kJ} \)[/tex]

4. Energy to vaporize water at [tex]\(100^\circ \text{C}\)[/tex]:

 Latent heat of vaporization [tex](\(\Delta H_{\text{vap}}\)) for water: \(40.67 \text{ kJ/mol}\).[/tex]

  Energy required to vaporize water:

[tex]\[ Q_4 = \text{moles of water} \times \Delta H_{\text{vap}} = 20.49 \text{ mol} \times 40.67 \text{ kJ/mol} = 833.34 \text{ kJ} \][/tex]

 [tex]\( Q_4 = 833.34 \text{ kJ} \)[/tex]

5. Energy to heat steam from [tex]\(100^\circ \text{C}\) to \(125^\circ \text{C}\):[/tex]

  Specific heat capacity of steam [tex](\(C_m\)): \(36.04 \text{ J/(mol} \cdot \text{°C)}\).[/tex]

  Energy required to heat steam:

[tex]\[ Q_5 = \text{moles of steam} \times C_m \times \Delta T = 20.49 \text{ mol} \times 36.04 \text{ J/(mol} \cdot \text{°C)} \times (125^\circ \text{C} - 100^\circ \text{C}) = 1831.79 \text{ J} \][/tex]

[tex]\( Q_5 = 1.83179 \text{ kJ} \)[/tex]

6. Total energy required.

  Summing up all the energy contributions.[tex]\[ Q_{\text{total}} = Q_1 + Q_2 + Q_3 + Q_4 + Q_5 = 7.48539 \text{ kJ} + 123.05 \text{ kJ} + 155.2978 \text{ kJ} + 833.34 \text{ kJ} + 1.83179 \text{ kJ} = 1120.00598 \text{ kJ} \][/tex]

Answer:

Explanation:

No reaction if there is no solid then the ions just stay there brah. All the ions remain in solution. We are all gonna make it!

Answer: a) 0.070 moles of oxygen were produced.

b) New pressure due to the oxygen gas is 2.4 atm

Explanation:

According to ideal gas equation:

[tex]PV=nRT[/tex]

P = pressure of gas = 2.7 atm

V = Volume of gas = 700 ml = 0.7 L

n = number of moles = ?

R = gas constant =[tex]0.0821Latm/Kmol[/tex]

T =temperature = 329 K

[tex]n=\frac{PV}{RT}[/tex]

[tex]n=\frac{2.7atm\times 0.7L}{0.0821 L atm/K mol\times 329K}=0.070moles[/tex]

Thus 0.070 moles of oxygen were produced.

When the 700 mL flask cools to a temperature of 293K.

[tex]PV=nRT[/tex]

[tex]P=\frac{nRT}{V}[/tex]

[tex]P=\frac{0.070\times 0.0821\times 293}{0.7}[/tex]

[tex]P=2.4atm[/tex]

The new pressure due to the oxygen gas is 2.4 atm

Answer:

what subject is this and how can u draw the structure :)

The reaction involves reduction of the given compound using NaBH4, followed by solvolysis with ethanol and protonation with H3O. Organic molecules structures are often simplified using the skeletal structure method. The stereochemistry of compounds can be represented using wedge and dash notation.

The given reaction involves the reduction of an organic compound using NaBH4, followed by the treatment with ethanol and then H3O. When a carbonyl compound reacts with NaBH4, the borohydride group of NaBH4 (it is a weak reducing agent with H-) adds to the carbonyl carbon. This process, known as reduction, transforms the carbonyl group to an alcohol. Following reduction, the compound is undergoing solvolysis in ethanol and then protonated with H3O. Due to the restriction of the platform, displaying the structural representation of the organic product is limited.

The wedge and dash notation can be used to represent the stereochemistry of the compounds. Solid wedges represent bonds coming up out of the plane whereas dashed lines represent bonds going into the plane. These techniques assist chemists to define molecular structures in three dimensions. Skeletal structure, a common way of simplifying complex organic structures in chemistry, represents carbons as each end of a line or bend in a line.

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Answer:

The phosphorylation of glucose to glucose-6-phosphate is endergonic reaction that is coupled to the exergonic hydrolysis of ATP.

Explanation:

In glycosis, the first reaction that takes place is the phosphorylation of glucose to glucose-6-phosphate by the enzyme hexokinase. This is an exergenic reaction. This is a coupled reaction in which phosphorylation of glucose is coupled to ATP hydrolysis. The free energy of ATP hydrolysis fuels glucose phosphorylation.

Answer:

The maximum volume is 122.73 mL

Explanation:

125 mL of butter that is 0.360 M in CH₃COOH and CH₃COO⁻.

pKa of acetic acid = 4.76

Using Henderson Hasselbalch equation.

pH =pKa + log [tex]\frac{CH_3COO^-}{CH_3COOH}[/tex]

= 4.76 + [tex]log \frac{0.3600}{0.3600}[/tex]

= 4.76

Assuming the pH of the buffer changes by a unit of 1 , then it will lose its' buffering capacity.

When HCl is added , it reacts with  CH₃COO⁻ to give CH₃COOH.

However, [CH₃COOH] increases and the log term  results in a negative value.

Let  assume, the new pH is less than 4.76

Let say 3.76; calculating the concentration when the pH is 3.76; we have

3.76 = 4.76 + log [tex]\frac{CH_3COO^-}{CH_3COOH}[/tex]

log [tex]\frac{CH_3COO^-}{CH_3COOH}[/tex] = 4.76 - 3.76

log [tex]\frac{CH_3COO^-}{CH_3COOH}[/tex] = 1.00

[tex]\frac{CH_3COO^-}{CH_3COOH}[/tex] = 0.1

Let number of  moles of acid be x   (i.e change in moles be x);  &

moles of acetic acid and conjugate base present be = Molarity × Volume

= 0.360 M × 125 mL

= 45 mmol

replacing initial concentrations and change in the above expression; we have;

[tex]\frac{[CH_3COO^-]}{[CH_3COOH]} = 0.1[/tex]

[tex]\frac{45-x}{45+x} =0.1[/tex]

0.1(45 + x) = 45 - x

4.5 + 0.1 x = 45 -x

0.1 x + x = 45 -4.5

1.1 x = 40.5

x = 40.5/1.1

x = 36.82

So, moles of acid added = 36.82 mmol

Molarity = 0.300 M

So, volume of acid = [tex]\frac{moles}{molarity }[/tex] =  [tex]\frac{36.82 \ mmol}{0.300}[/tex]

= 122.73 mL

Answer:

[tex]\large \boxed{\text{103 kPa}}[/tex]

Explanation:

We can use the Ideal Gas Law — pV = nRT

Data:

V = 66.8 L

m = 77.8 g

T = 25 °C

Calculations:

(a) Moles of N₂

[tex]\text{Moles of N}_{2} = \text{77.8 g N}_{2} \times \dfrac{\text{1 mol N}_{2}}{\text{28.01 g N}_{2}} = \text{2.778 mol N}_{2}[/tex]

(b) Convert the temperature to kelvins

T = (25 + 273.15) K = 298.15 K

(c) Calculate the pressure

[tex]\begin{array}{rcl}pV & =& nRT\\p \times \text{66.8 L} & = & \text{2.778 mol} \times \text{8.314 kPa$\cdot$ L$\cdot$K$^{-1}$mol$^{-1}\times$ 298.15 K}\\66.8p & = & \text{6886 kPa}\\p & = & \textbf{103 kPa}\end{array}\\\text{The pressure in the bag is $\large \boxed{\textbf{103 kPa}}$}[/tex]

Answer:

2-Tert-butyl-cyclohexanone, 2,6-Di-Tert-Butylcyclohexanone and Di-isopropyl-amine.  

Explanation:

The balanced chemical equation for the overall chemical reaction is 2NO2(g) + F2(g) → 2NO2F(g). The experimentally observable rate law for the overall chemical reaction is Rate = k [NO2]^2 [F2]. The rate constant k for the overall chemical reaction can be expressed in terms of k1, k2, k-1, and k-2 as k = k1 (k2/k-1).

The balanced chemical equation for the overall chemical reaction is:

2NO2(g) + F2(g) → 2NO2F(g)

The experimentally observable rate law for the overall chemical reaction is:

Rate = k [NO2]2 [F2]

The rate constant k for the overall chemical reaction can be expressed in terms of k1, k2, k-1, and k-2 as:

k = k1 (k2/k-1)

Answer:

celulares por que tienen una fuente de energiq

Answer:

[tex]T = 34.493\,^{\textdegree}C[/tex]

Explanation:

The equilibrium temperature of the gas-container system is:

[tex]Q_{Cu} = -Q_{g}[/tex]

[tex](0.3\,kg)\cdot\left(386\,\frac{J}{kg\cdot ^{\textdegree}C} \right)\cdot (T-20\,^{\textdegree}C) = (1.5\,mol)\cdot \left(12.5\,\frac{J}{mol\cdot ^{\textdegree}C} \right)\cdot (124\,^{\textdegree}C-T)[/tex]

[tex]\left(115.8\,\frac{J}{^{\textdegree}C} \right)\cdot (T-20\,^{\textdegree}C) = \left(18.75\,\frac{J}{^{\textdegree}C} \right)\cdot (124\,^{\textdegree}C-T)[/tex]

[tex]134.55\cdot T = 4641[/tex]

[tex]T = 34.493\,^{\textdegree}C[/tex]

The most likely substrates for the oligo-(α1→4)-(α1→4) glucanotransferase catalytic center of glycogen debranching enzyme are maltosyl and maltotriosyl residues.

The glycogen debranching enzyme has two activities; it acts as a glucosidase and a glucanotransferase. The (α‑6) glucosidase activity hydrolyzes α-1,6 glycosidic bonds, whereas the oligo-(α1→4)-(α1→4) glucanotransferase activity shifts α-1,4-linked glucose chains from one branch to another, often creating a substrate that the glucosidase can act upon. Based on the mechanism of glycogenolysis, the enzyme glycogen phosphorylase releases glucose units from the linear chain until a few are left near the branching point, and it is here that the glucan transferase action is relevant. The glucan transferase shifts the remaining α-1,4 linked glucose units, leaving a single α-1,6 linked glucose that the glucosidase can then release. Hence, the most likely substrates for the oligo-(α1→4)-(α1→4) glucanotransferase catalytic center are maltosyl and maltotriosyl residues, as these are the short α-1,4 linked glucose chains that are left after phosphorylase action.

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Answer:

742.2 L

Explanation:

First we must find the number of moles of nitroglycerine reacted.

Molar mass of nitroglycerine= 227.0865 g/mol

Mass of nitroglycerine involved = 1×10^3 g

Number of moles of nitroglycerine= 1×10^3g/227.0865 g/mol

n= 4.40361 moles

T= 1985°C + 273= 2258K

P= 1.100atm

R= 0.082atmLmol-1K-1

Using the ideal gas equation:

PV= nRT

V= nRT/P

V= 4.40361× 0.082× 2258/1.1

V= 742 L

Considering the reaction stoichiometry and ideal gas law, the volume of gas produced is 5375.626 L.

The balanced reaction is:

4 C₃H₅N₃O₉(s) → 12 CO₂(g) + 10 H₂O(g) + 6 N₂(g) + O₂(g)

By reaction stoichiometry (that is, the relationship between the amount of reagents and products in a chemical reaction), the following amounts of moles of each compound participate in the reaction:

Then 4 moles of nitroglycerine C₃H₅N₃O₉(s) produce in total 29 moles of gas [12 moles of CO₂(g) + 10 moles of H₂O(g) + 6 moles of N₂(g) + 1 mole O₂(g)]

Being the molar mass of nitroglycerine C₃H₅N₃O₉(s) 227 g/mole, then the amount of moles that 1 kg (1000 g) of the compound contains can be calculated as:

[tex]1000 gramsx\frac{1 mole}{227 grams}= 4.405 moles[/tex]

Then you can apply the following rule of three: if by stoichiometry 4 moles of nitroglycerine C₃H₅N₃O₉(s) produce in total 29 moles of gas, 4.405 moles of C₃H₅N₃O₉(s) will produce how many moles of gas?

[tex]amount of moles of gas=\frac{4.405 moles of nitroglycerinex29 moles of gas}{4 moles of nitroglycerine}[/tex]

amount of moles of gas= 31.93625 moles

On the other hand, an ideal gas is characterized by three state variables: absolute pressure (P), volume (V), and absolute temperature (T). The relationship between them constitutes the ideal gas law, an equation that relates the three variables if the amount of substance, number of moles n, remains constant and where R is the molar constant of the gases:

P× V = n× R× T

In this case, you know:

Replacing:

1.1 atm× V= 31.93625 moles× 0.082 [tex]\frac{atmL}{molK}[/tex]× 2258 K

Solving:

V= (31.93625 moles× 0.082 [tex]\frac{atmL}{molK}[/tex]× 2258 K) ÷ 1.1 atm

V= 5375.625 L

Finally, the volume of gas produced is 5375.626 L.

Learn more:

Answer:

Explanation:

check the attachment below for correct explanations.

Answer:

Moles = 0.375

Explanation:

Moles= m/M

= 12/32 = 0.375mol

There are approximately 0.375 moles of O₂ in 12.0 grams of O₂.

To find the number of moles in a given mass of a substance, one can use the formula:

[tex]\[ \text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}} \][/tex]

The molar mass of oxygen (O₂) is approximately 32.0 g/mol, since one oxygen atom has an atomic mass of approximately 16.0 g/mol and O₂ has two oxygen atoms.

Given the mass of O₂ is 12.0 g, we can calculate the moles as follows:

[tex]\[ \text{moles of \ } {O_2} = \frac{12.0 \text{ g}}{32.0 \text{ g/mol}} \][/tex]

[tex]\[ \text{moles of} \ {O_2} = 0.375 \text{ mol} \][/tex]

Therefore, there are 0.375 moles of O₂ in 12.0 grams of O₂.

Answer:

Explanation:

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To construct a simulated 1H NMR spectrum for 1-chloropropane, you need to consider the chemical shifts and splitting patterns of the hydrogen atoms. Follow these steps: identify the different types of hydrogen atoms, determine the chemical shifts, assign integration values, and drag and drop splitting patterns.

To construct a simulated 1H NMR spectrum for 1-chloropropane, you need to consider the chemical shifts and splitting patterns of the hydrogen atoms. Here's a step-by-step guide:

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The final pressure of the ammonia gas, given that the the volume decreased by 50.0%, is 11 atm

How to calculate the final pressure f the ammonia gas?

To solve this question in the most simplified way, we shall begin by listing out the given data from the question. This is shown below:

The final pressure of the ammonia gas can be calculated as shown below:

[tex]\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}\\\\\frac{4.6\ \times\ V}{252} = \frac{P_2\ \times\ 0.5V}{296}\\\\252\ \times\ P_2\ \times\ 0.5V = 4.6\ \times\ V\ \times\ 296\\\\P_2 = \frac{4.6\ \times\ V\ \times\ 296}{252\ \times\ 0.5V} \\\\P_2 = 11\ atm[/tex]

Thus, the final pressure is 11 atm

Answer:

15.1 seconds is the half life of the reaction when concentration of the substrate is 2.77 M.

Explanation:

A → B + C

The rate law of the reaction will be :

[tex]R=k[A]^x[/tex]

Initial rate of the reaction when concentration of the substrate was 0.4 M:

[tex]0.183 M/s=k[0.4 M]^x[/tex]..[1]

Initial rate of the reaction when concentration of the substrate was 0.8 M:

[tex]0.183 M/s=k[0.8 M]^x[/tex]...[2]

[1] ÷ [2] :

[tex]\frac{0.183 M/s}{0.183 M/s}=\frac{k[0.4 M]^x}{k[0.8 M]^x}[/tex]

x = 0

The order of the reaction is zero.

For the value of rate constant ,k:

[tex]0.183 M/s=k[0.4 M]^x[/tex]..[1]

x = 0

[tex]0.183 M/s=k[0.4 M]^0[/tex]

k= 0.183 M/s

The half life of the zero order kinetics is given by :

[tex]t_{1/2}=\frac{[A_o]}{2k}[/tex]

Where:

[tex][A_o][/tex] = Initial concentration of A

k = Rate constant of the reaction

So, the half-life for the decomposition of substrate 1 when the initial concentration of the substrate is 2.77 M:

[tex]t_{1/2}=\frac{2.77 M}{0.183 M/s}=15.1 s[/tex]

15.1 seconds is the half life of the reaction when concentration of the substrate is 2.77 M.

The half-life for the decomposition of substrate 1 when the initial concentration is 2.77 M is approximately 18 minutes.

The half-life of a reaction is the time required for one-half of a given amount of reactant to be consumed. In this case, we have initial rate data for the decomposition of substrate 1 with varying concentrations. To determine the half-life of substrate 1, we need to find the time it takes for the concentration to decrease to half its initial value.

Based on the given data, we can see that the initial rate of the reaction is constant at 0.183 M/s for different initial concentrations of substrate 1 (0.4 M, 0.8 M, and 2 M). Since the initial concentration of substrate 1 is given as 2.77 M, it will take the same amount of time as the other initial concentrations to reach half its initial concentration.

Therefore, the half-life for the decomposition of substrate 1 when the initial concentration of the substrate is 2.77 M is the same as the half-life observed when the initial concentration is 0.4 M, 0.8 M, or 2 M, which is approximately 18 minutes.

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Answer:

pH  = 13.09

Explanation:

Zn(OH)2 --> Zn+2 + 2OH-   Ksp = 3X10^-15

Zn+2 + 4OH-   --> Zn(OH)4-2   Kf = 2X10^15

K = Ksp X Kf

  = 3*2*10^-15 * 10^15

  = 6

Concentration of OH⁻ = 2[Ba(OH)₂] = 2 * 0.15 = 3 M

                Zn(OH)₂ + 2OH⁻(aq)  --> Zn(OH)₄²⁻(aq)

Initial:           0             0.3                      0

Change:                      -2x                     +x

Equilibrium:               0.3 - 2x                 x

K = Zn(OH)₄²⁻/[OH⁻]²

6 = x/(0.3 - 2x)²  

6 = x/(0.3 -2x)(0.3 -2x)

6(0.09 -1.2x + 4x²) = x

0.54 - 7.2x + 24x² = x

24x² - 8.2x + 0.54 = 0

Upon solving as quadratic equation, we obtain;

x = 0.089

Therefore,

Concentration of (OH⁻) = 0.3 - 2x

                                    = 0.3 -(2*0.089)

                                  = 0.122

pOH = -log[OH⁻]

         = -log 0.122

          = 0.91

pH = 14-0.91

     = 13.09

Answer:

B

Explanation:

B. There are two atoms of Nitrogen and two atoms of Hydrogen combined to make Ammonia.

Answer:

The specific heat of copper is  [tex]C= 392 J/kg\cdot ^o K[/tex]

Explanation:

From the question we are told that

The amount of energy contributed by each oscillating lattice site  is  [tex]E =3 kT[/tex]

       The atomic mass of copper  is  [tex]M = 63.6 g/mol[/tex]

        The atomic mass of aluminum is  [tex]m_a = 27.0g/mol[/tex]

        The specific heat of aluminum is  [tex]c_a = 900 J/kg-K[/tex]

 The objective of this solution is to obtain the specific heat of copper

       Now specific heat can be  defined as the heat required to raise the temperature of  1 kg of a substance by  [tex]1 ^o K[/tex]

  The general equation for specific heat is  

                    [tex]C = \frac{dU}{dT}[/tex]

Where [tex]dT[/tex] is the change in temperature

             [tex]dU[/tex] is the change in internal energy

The internal energy is mathematically evaluated as

                       [tex]U = 3nk_BT[/tex]

      Where  [tex]k_B[/tex] is the Boltzmann constant with a value of [tex]1.38*10^{-23} kg \cdot m^2 /s^2 \cdot ^o K[/tex]

                    T is the room temperature

                      n is the number of atoms in a substance

Generally number of  atoms in mass of an element can be obtained using the mathematical operation

                      [tex]n = \frac{m}{M} * N_A[/tex]

Where [tex]N_A[/tex] is the Avogadro's number with a constant value of  [tex]6.022*10^{23} / mol[/tex]

          M is the atomic mass of the element

           m actual mass of the element

  So the number of atoms in 1 kg of copper is evaluated as  

             [tex]m = 1 kg = 1 kg * \frac{10000 g}{1kg } = 1000g[/tex]

The number of atom is  

                       [tex]n = \frac{1000}{63.6} * (6.0*0^{23})[/tex]

                          [tex]= 9.46*10^{24} \ atoms[/tex]

Now substituting the equation for internal energy into the equation for specific heat

          [tex]C = \frac{d}{dT} (3 n k_B T)[/tex]

              [tex]=3nk_B[/tex]

Substituting values

         [tex]C = 3 (9.46*10^{24} )(1.38 *10^{-23})[/tex]

            [tex]C= 392 J/kg\cdot ^o K[/tex]

6.81% is the percent yield when 12 g of CO2 are formed experimentally from the reaction of 0.5 moles of C8H18 reacting with excess oxygen.

Explanation:

Data given:

mass of carbon dioxide formed = 12 grams (actual yield)

atomic mass of CO2 = 44.01 grams/mole

moles of [tex]C_{8} H_{18}[/tex] = 0.5

Balanced chemical reaction:

2 [tex]C_{8} H_{18}[/tex] + 25 [tex]O_{2}[/tex] ⇒ 16 C[tex]O_{2}[/tex] + 18 [tex]H_{2} _O{}[/tex]

number of moles of carbon dioxide given is

number of moles = [tex]\frac{mass}{atomic mass of 1 mole}[/tex]

number of moles= [tex]\frac{12}{44}[/tex]

number of moles of carbon dioxide gas = 0.27 moles

from the reaction 2 moles of [tex]C_{8} H_{18}[/tex] reacts to produce 16 moles of C[tex]O_{2}[/tex]

So, when 0.5 moles reacted it produces x moles

[tex]\frac{16}{2}[/tex] = [tex]\frac{x}{0.5}[/tex]

x = 4

4 moles of carbon dioxide formed, mass from it will give theoretical yield.

mass  = number of moles x molar mass

mass = 4 x 44.01

        = 176.04 grams

percent yield =[tex]\frac{actual yield}{theoretical yield}[/tex] x 100

percent yield = [tex]\frac{12}{176.04}[/tex] x 100

     percent yield  = 6.81 %

n = 0.0989 moles

n = PV / RT

P = 2.09atm

V = 1.13L

R = 0.08206

T = 291K

Plug the numbers in the equation.

n = (2.09atm)(1.13L) / (0.08206)(291K)

n = 0.0989 moles

Final answer:

Using the ideal gas law PV = nRT, and rearranging for n (moles), we plug in the given values of pressure, volume, and temperature alongside the ideal gas constant to calculate the number of moles of gas present. There are approximately 0.0988 moles of gas present in 1.13 L of gas at [tex]\(2.09 \, \text{atm}\) and \(291 \, \text{K}\).[/tex]

Explanation:

To calculate the number of moles of gas present, we can use the ideal gas law equation:

[tex]\[ PV = nRT \][/tex]

First, we need to convert the given pressure to atm and the volume to liters if they are not already in those units.

Given:

- [tex]\( P = 2.09 \, \text{atm} \)[/tex]

-[tex]\( V = 1.13 \, \text{L} \[/tex]

- [tex]\( T = 291 \, \text{K} \)[/tex]

Now, we can rearrange the ideal gas law equation to solve for \( n \):

[tex]\[ n = \frac{PV}{RT} \][/tex]

[tex]\[ n = \frac{(2.09 \, \text{atm})(1.13 \, \text{L})}{(0.0821 \, \text{atm} \cdot \text{L} / \text{mol} \cdot \text{K})(291 \, \text{K})} \][/tex]

[tex]\[ n = \frac{2.3617}{23.9211} \][/tex]

[tex]\[ n \approx 0.0988 \, \text{mol} \][/tex]

Therefore, There are approximately 0.0988 moles of gas present in 1.13 L of gas at [tex]\(2.09 \, \text{atm}\) and \(291 \, \text{K}\).[/tex]